honors instruction manual - tenafly public schools

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Honors Chemistry Lab:

Lab Manual®

Tenafly High School

©September 2014

2nd edition

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Lab # CHEMISTRY LAB - ACTIVITY TITLES Instruction

Manual Page

- Preparing for Lab 3

- Performing the Lab 4

- Writing the Lab Report 5

- Grading Template 7

1 Safety Lab 9

2 Percent Composition of a Mixture 13

3 Comparing Physical and Chemical Changes 14

4 Determining an Empirical Formula 17

5 Composition of Hydrates 19

6 Determining the Thickness of Aluminum Foil 21

7 Types of Chemical Reactions 22

8 Relating Moles to Coefficients in an Equation 26

9 Particle Size from Collision Probabilities 28

10 Flame Tests and Emission Spectroscopy 30

11 Investigation of the Hydrogen Spectrum 32

12 Determining the Half-life of Ba-137m 34

13 Periodic Table: A Study of the Reactivity of Metals 36

14 Preparation and Properties of Oxygen 38

15 Model Construction: Polarity and Shape 41

16 Paper Chromatography 43

17 Determination of Absolute Zero 45

18 Molar Volume of a Gas 47

19 Construction of an Air Bag 50

20 A Study of Phase Change Using Lauric Acid 52

21 Application of Solubility Rules 55

22 The 6-Solution Puzzle 57

23 Rate of a Chemical Reaction 59

24 A Study of Equilibrium 62

25 Change in Enthalpy of a Reaction 65

26 Determination of pH Using Indicators 67 (in progress)

27 Acid-Base Titration

28 Decomposition of Baking Soda

29 The Ice Cream Lab

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PREPARING FOR LAB

Before actually performing an experiment, you need to know WHAT you are doing and, more

importantly, WHY you are doing what you are doing!

PRELAB (30 pts): Before each lab, you will be responsible for writing the following in your lab

notebook (a black and white composition notebook).

• You will NOT be allowed to take anything but your lab notebook into the lab area, and you will NOT have

access to a lab manual during the lab – your level of preparation will determine your success in the

laboratory! If your prelab is not completed, you must complete it for half credit BEFORE being allowed to

perform the lab.

• You must independently complete your prelab in your own handwriting. It should not be photocopied or

typed/ printed on a computer.

PURPOSE (6 pts): What question/ hypothesis does this experiment seek to answer/ test? Read

through the lab manual and procedures for the lab, and come up with the specific question or

hypothesis that this experiment can answer or test.

BACKGROUND RESEARCH (6 pts): What do you need to know before doing this experiment? You will

be given terms to define or questions to answer specific to your experiment. The answers may be

found by researching online, looking in your book or in your class notes, or in the lab background.

MATERIALS (6 pts): What do you need to perform this experiment, and how much? List all equipment

and reagents (with concentrations for solutions). Subdivide materials into the parts if there are

different parts to the lab.

PROCEDURES & OBSERVATIONS/ DATA (6 pts): What will you do and what tests will you perform to

address your purpose? Split your page in half. On the left, write your procedures. On the right, leave

space for observations and notes, and create data tables for the measurements that you will make.

� (3) PROCEDURES: Summarize what you are doing in this experiment—give enough detail so that

you can perform the lab without the manual. Use your own words or feel free to draw diagrams

instead of words.

� (3) DATA TABLES: During the lab, all data taken during the experiment must be written in ink and

identified (i.e. a random mass without any indication what was massed is fairly useless), with the

correct metric unit for all measurements. Nothing is to be written on scrap paper and transferred

later. If you make an error or repeat the procedure, draw a single line through the data and write

the corrected data next to it. There should be no erasing or white-out. I will initial the data at the

end of the lab. Each person in a group must have a complete set of data recorded before leaving

the lab—you may not have one person record the data, and share the data later.

SAFETY CONCERNS (6 pts): With what chemicals and procedures must I exercise extra caution? How

can reduce the risk? Whenever there is a hazardous substance (e.g. corrosive, stains skin, flammable)

or procedural hazard, please write the substance or physical hazard, associated danger, as well as the

preventative measure taken to reduce risk.

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PERFORMING THE LAB

TECHNIQUE, SAFETY, CLEANUP, GROUP DYNAMICS (10 pts): During the lab, you

will be assessed on your ability to perform the lab safely and efficiently. Inappropriate behavior (e.g.

sitting on desks, horseplay, discarding waste improperly, careless placement of glassware), inequitable

division of labor, poor group communication, and incomplete cleanup will affect your grade here.

Individuals in the same group may be assessed different grades. For a group of four, please divide

yourselves into the following roles so that responsibilities for each person are clear. For a group of

three, one person will be both the workflow manager and the materials manager. For a group of two,

one person will be both the workflow manager and technician, and the other person will be the

materials manager and the quality control manager.

I. WORKFLOW MANAGER

� Keeps group on task and out of trouble—makes sure that the group finishes the lab on time!

� Helps and fills in as needed to make lab run smoothly—should NEVER be idly doing nothing!

� Answers group questions with common sense and own resources before approaching teacher

� Only member who addresses the teacher (unless there is an emergency) to bring concerns,

problems, or questions to the teacher (unless there is an emergency).

II. MATERIALS MANAGER

� Acquires lab aprons for groups and materials (both chemicals and glassware/ equipment)—makes

sure that all members are familiar with names and properties of important materials!

� Sets up lab apparatuses and makes sure glassware is clean (and dry, for certain labs)

� Oversees proper safety precautions/ practices (e.g., notifying teacher of chemical spills, tying hair

back when working with fire, using insulated gloves when handling hot objects)

� Cleans up, or if there is a lot of cleanup, equitably distribute cleanup duties

III. TECHNICIAN

� Reads at least 2 steps ahead before performing a lab step.

� Performs the procedure or oversees/ delegates steps to other members if a multi-part lab—makes

sure that all members understand the procedures and measurements!

� Signs initials in all lab notebooks next to each major step of the procedure immediately after that

step has been completed.

� Makes and labels/ records measurements to proper number of significant figures with proper units.

IV. QUALITY CONTROL MANAGER

� Ensures procedures are done properly, assisting technician to ensure that measurements and

observations are recorded neatly and labeled correctly for the appropriate steps.

� Also signs initials next to each major step of the procedure after that step has been completed.

� Assesses whether the data is consistent and logical, and if not, makes note of any errors that may

affect the data—If the data seem to have significant error, he/ she is responsible for bringing it to

the attention of the workflow manager asap and assessing whether the error is significant

enough to warrant a repeat or a request for sample lab data.

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WRITING THE LAB REPORT

Whenever scientific work is done, work be clearly documented and published in a peer-reviewed

journal so that other scientists can reproduce experiments and independently verify any conclusions.

Therefore, when you do a lab, you will learn to document and publish your work in the form of a

formal laboratory report. Lab reports will be written collaboratively (at least at first!), with lab

members cycling through writing each part or editing the overall report for clarity, formatting,

grammar, and spelling.

LAB REPORT FORMAT/ SPELLING/ GRAMMAR (15 pts): Your formatting (i.e. correct order,

including all parts, spacing, providing titles for each part), spelling, and grammar are assessed

PART 1: INTRODUCTION & PROCEDURE (15 pts) I. HEADING:

a) NAMES & ROLES: Include name of all group members—first and last names, spelled correctly.

Include each group member’s portion of the lab report written after the name in parentheses.

Scientists give credit to their collaborators when a work is published—so should you!

b) COURSE TITLE/ INSTRUCTOR: Mrs. Woleslagle, Period X Honors Chemistry

c) DATE DUE: Please include the date that the lab is being handed in (not the day that the lab was

performed) in the Month/Day/ Year

d) TITLE: Please include the title given in the lab manual, centered and underlined.

II. INTRODUCTION: The introduction must be written using complete sentences in your own words

(i.e. do not copy directly from the lab handbook, textbook, or lab partner).

a) State the PURPOSE or GOALS of the experiment. If a hypothesis is to be tested in the

experiment, it must be written as an “If… then…” statement.

b) Give a detailed discussion of the BACKGROUND needed to understand the results. This section

needs to present the background information, previous research, and/ or formulas that the

reader needs to understand and solve the problem. For some labs, you will provided with a list of

key terms/ concepts to be included.

IV. MATERIALS: Scientific publications always include lists of the reagents and the equipment used, so

that other scientists trying to repeat or extend the experiments can carry them out under similar

conditions. You need to include all reagents used, as well as the lab equipment used. As an option,

you may include a labeled drawing or photograph of your setup.

V. PROCEDURES: Scientists include a description of the actual procedure used in their reports, so that

other scientists can replicate the experiment if needed. You should include a numbered list, with

the steps that you actually performed in the correct sequence, described in your own words (not

copied from your lab partner). The procedure should include enough detail so that you could give

your procedure to a classmate and he/ she could correctly carry it out. It is not acceptable to copy

the procedure word for word from the lab handout. You do not need to include safety equipment

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(goggles, aprons) in your procedure. The procedures must be written in passive voice and past

tense (see the course website for more information).

PART 2: DATA & CALCULATIONS (15 pts) All data used or collected for your experiment must be presented in an appropriate format such as a

table or graph.

I. TABLES: Each table must have a descriptive title, and all measurement must be reported with the

correct metric unit. Any measured data must be put into a table, and any calculated data must be

put into a separate table. Percentage error should be put in the “calculated data” table.

GRAPHS: Each graph must have a descriptive title, with both axes labeled and proper units

specified. Graphs may be computer generated and done on graph paper and staples or taped

inside the report.

II. CALCULATIONS: If calculations are required, state the formulas used and show sample calculations

(you do not need to do every single calculation – just a sample). Each calculation should appear on

a separate line. Make sure all of your work is shown, measurements have proper units, and

significant figure rules are followed for all measurements and calculations. Percent error

calculations should be shown here. You may NEATLY handwrite calculations.

PART 3: QUESTIONS & CONCLUSIONS (15 pts)

I. ANALYZE & APPLY QUESTIONS: Scientists include a discussion of the significance of the work and

applications to different areas. This is your opportunity to show me (your teacher) your

understanding of the concepts and to extend them to other situations. Show me what you can do!

Number the questions, skip a line between questions, and answer in complete sentences with

grammatically correct English. If the question is a math-based problem, show all your work and

include all relevant formulas, significant figures, etc. You may NEATLY handwrite calculations.

II. CONCLUSIONS: The conclusion refers back to the purpose, summarizes the results, and indicates

any changes you would make to the experiment if you were to repeat it. The conclusion should be

in paragraph style and it should be brief. This is not a place to give your personal opinion about the

enjoyment value of the lab—it should be a critical assessment of the results of the experiment.

a) Summarize your RESULTS

CLAIMS—What can you claim? Have you accomplished your purpose? What are the results of

your experiment (quantitative and/or qualitative)? What are your claims or conclusions

after having performed this lab? If there is an unknown, the number of the unknown must

be included in the identification.

EVIDENCE—How do you know? Why are you making these claims? What evidence backs up

your claims/ conclusions? Refer back to your data for support.

b) Discuss EXPERIMENTAL and/ or PROCEDURAL SOURCES OF ERROR.

What portions of the lab should you do differently/ redesign to avoid or reduce these errors?

Why? What portions of the lab should you repeat? Why? Describe how these sources of error

affected your quantitative results. Did they make your results too big? Too small? Why?

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Never include “human error” or mistakes in measurement or calculations as a sources of

error—stupidity or negligence are never acceptable errors.

ABSENCES/ PRELABS. If you are absent the day before a lab, it is your responsibility to find out

what lab will be done and write your pre-lab, since it is included as part of the grade. If you do not

have the prelab done at the beginning of the lab, the grade on that part is ZERO. If you are absent on

the day of a lab, it is your responsibility:

(i) to show the prelab to your teacher on the day of your return;

(ii) to obtain the lab data from your partner(s) on the following day, and:

(iii) to arrange a makeup date for the performance of the lab/ a due date for the lab report

with your teacher.

GRADING. The report must be typed and will be due the following day at the beginning of class.

When labs were done as partners or groups, everyone must submit a completed report. The reports

from partners should be attached and handed in as a packet. Sometimes all reports will be graded

individually; at other times, one will be selected at random to represent the group. Final grades for

members of the team may still vary because of the prelab, lab performance, organization, etc. Points

will be deducted for reports handed in after the due date. Graded reports must be saved in your binders. They will be collected as a packet at the end of

each quarter and graded for completeness. Be prepared for lab quizzes and/ or questions pertaining to

labs on tests.

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GRADING TEMPLATE

NAME: __________________________________ LAB #:____________ DATE: ______________

SECTION YOUR

SCORE

MAX

SCORE

PRELAB

• Purpose (6 pts)

• Background Research (6 pts)

• Materials (6 pts)

• Procedures/ Data Tables (6 pts)

• Safety Concerns (6 pts)

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LAB PERFORMANCE

• Workflow manager –keeps group on task, makes sure lab is finished

on time, sole teacher contact

• Materials manager –sets up & oversees cleanup, manages safety

• Technician –performs lab, ensures measurements made with correct

number of sig figs, signs-off on procedures

• Quality Control (QC) manager –assists technician, double checks

and signs-off on procedure, ensures proper recording of measurements,

manages error

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LAB REPORT

• Formatting, Spelling, & Grammar

• Part 1: Introduction, Materials, Procedures

• Part 2: Data Tables, Graphs, Calculations

• Part 3: Questions, Conclusion, Error Analysis

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TOTAL LAB SCORE 100

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1. SAFETY LAB

INTRODUCTION Welcome to Honors Chemistry. One of the most important considerations during any

laboratory activity is safety. Labs provide a valuable learning experience (and can be fun as well) as

long as common sense and safe behavior are practiced. This activity will help you establish skills you

will use all year—your application of safety rules and your recognition of frequently used lab

equipment. The Flinn safety contract and lab equipment diagrams at the front of your lab manual are

important resources for this lab.

BACKGROUND RESEARCH 1. List 3-4 general protective/ preventative measures YOU should take to minimize your exposure to

hazardous chemicals.

2. There are a few common hazardous chemical categories that you will encounter throughout this

year.

-Define the following categories.

-Give one example (chemical name and formula) for each category.

a. Corrosive

b. Oxidizer

c. Volatile

d. Alkali

3. What lab equipment is best for measuring volumes precisely? (Hint: it comes in different sizes).

4. Look up the Safety Data Sheet (SDS) for hydrochloric acid. What is its CAS number, and what are

the hazards associated with this substance?

REVIEW YOUR PICTURES OF LAB EQUIPMENT.

SAFETY CONCERNS Hydrochloric acid (HCl) is hazardous – please refer to the SDS. Always wear goggles, gloves, and

aprons, and flush with water for at least 5 minutes if skin contact occurs.

MATERIALS REAGENTS EQUIPMENT

Nylon

Albumin

6 or 10 M

hydrochloric acid

(HCl)

watch glass

forceps

pipet

beaker

wash bottle

Bunsen burner

striker

numbered lab

equipment

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PROCEDURES

Overview: Start at the station where you are seated, and stay with your lab table group. Perform the activity as

stated and record information on the data table. Clean up after you are finished with each station.

When entire group is ready, go to the another station – stations can be done in any order. When

finished with all stations, return to your original station, wipe down table and clean up (do not remove

goggles and aprons during cleanup). Remove goggles and aprons only right before returning to

classroom desks.

#1: Effect of Acid on CLOTHING Use the forceps to take one piece of nylon from the beaker and put it on the watch glass. Use the

small dropper bottle filled with ACID to drop 2-3 drops of acid on the nylon. OBSERVE ANY CHANGES

TO THE NYLON.

� CLEANUP: Use the wash bottle to rinse the nylon and acid off of the watch glass into the waste

beaker. Wipe watch glass clean and dry with paper towel, and return to station. Put paper towel in

garbage.

#2: Effect of Acid on SKIN or EYES Use the plastic pipet to take 2-3 drops of albumin (egg white) from the beaker and put it on the watch

glass. Use the small dropper bottle filled with ACID to drop 2-3 drops of acid on the albumin. OBSERVE

ANY CHANGES TO THE ALBUMIN.

� CLEANUP: Rinse the albumin and acid off of the watch glass into the waste beaker. Wipe watch

glass clean and dry with paper towel, and return to station. Put paper towel in garbage.

#3: NOT ready to begin lab work Imagine your lab group is set up to work at this lab table, and the equipment you need for this lab has

been placed at your table. DO NOT CHANGE THE SETUP OF THIS TABLE – YOU ARE ONLY OBSERVING.

Draw a quick sketch of the table, or if you are a terrible artist – take a picture with your cell phone,

which you can print out and paste into your lab notebook. Observe/ list at least 4 things that need to

be fixed/ addressed BEFORE you can begin lab work.

#4: Lighting a BUNSEN BURNER

Watch the instructor demonstrate the technique of how to light the Bunsen burner with the striker.

Try it yourself. Sketch the burner, or if you are a terrible artist—take a picture with your cell phone,

which you can print out and paste into your lab notebook. Make sure to label the barrel, needle valve,

base, and hose. Make a note of what the gas valve looks like relative to the hose when the gas is on

and when it is off (perpendicular or parallel).

#5: Safety equipment DO NOT CHANGE THE SETUP OF THIS TABLE – YOU ARE ONLY OBSERVING. Identify the two important pieces

of safety equipment located at this station and when you would use them.

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Station #6: Equipment & Use DO NOT CHANGE THE SETUP OF THIS TABLE – YOU ARE ONLY OBSERVING. Identify each numbered piece of lab

equipment AND give a brief description of its use. You should sketch or take a picture of each piece so that you

can recognize it in the future.

DATA 1. In a “before” and “after” data table, summarize your observations from your experiment at station

1 and 2 about the effect of acid on clothing (i.e. nylon) and eye/ body tissue (i.e. albumin).

2. Either resketch/ scan (from your lab notebook) and paste in a picture of station 3.

3. Either resketch/ or scan (from your lab notebook) and paste in a picture of the burner from station

4, and (only) label the barrel, needle valve, base, and hose.

4. In a data table, include the number, name, picture/ sketch, and function of each piece of

equipment from station 6.

(You do not need to include a table or picture from station 5)

ANALYZE & APPLY QUESTIONS

1. Based on what you saw at station 1 and 2, what is the effect of acid on clothing and eye/ body

tissue, and what would you do in the lab to prevent exposure to your clothing and eye/ body

tissue?

2. In station 3, what are the four things that must be fixed before starting the experiment?

3. In station 4,

a. What part of the Bunsen burner controls flame color? How do you get a yellow flame? A

blue flame? Which color flame is a clean flame?

b. What part of the Bunsen burner controls flame height? How do you get a big flame? A

small flame? What part of the flame is the hottest part?

c. How do you turn the Bunsen burner off or on?

4. What are the two pieces of safety equipment located at station 5?

CONCLUSION

Complete a standard conclusion and error analysis – see “Writing the Lab Report” for help.

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2. PERCENT COMPOSITION OF A MIXTURE

INTRODUCTION In a mixture, whether homogeneous or heterogeneous, the components may be present in

any proportion. In order to determine the percent composition, the components must be separated

by some means, usually physical. All physical methods of separation work by exploiting differences in

physical properties. For example, crystallization uses differences in solubility at different

temperatures to crystallize a solid out of solution (i.e., a hot saturated solution is cooled, and the least

soluble solid precipitates first).

In this lab, you will think about the physical properties of the various components making up

your mixture, and then determine procedures that will help you separate the parts.

PURPOSE Determine the % composition of a mixture of sand (silicon dioxide, SiO2), table salt (sodium chloride,

NaCl), and iron filings (Fe).

BACKGROUND RESEARCH 1. Define the following methods of separation and give an example of a mixture that could be

separated using the method:

a. Filtration

b. Distillation

c. Evaporation

d. Chromatography

e. Magnetic Separation

2. In a mixture made of component A and B, what is the formula for the percent by mass of

component A?

SAFETY CONCERNS Wear your goggles and aprons! You do not need to wear gloves for this lab. Depending on your

procedures, you may need to observe other safety precautions.

MATERIALS REAGENTS EQUIPMENT

Sample of mixture (sand, salt, iron filings =

SiO2, NaCl, Fe)

Any equipment needed to carry out the

procedure

PROCEDURES Working in teams, design your experiment, design your experiment:

• Write your procedures in list format—

include cleanup!!! The three (dry)

components of your mixtures should be

submitted at the end of the lab.

• Identify safety precautions.

• Write down your equipment

• Prepare a data table

Submit your procedure for approval. Then, carry out your experiment, record data, and perform the

necessary calculations.

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DATA & CALCULATIONS 1. Make a data table for any measurements or calculated numbers.

2. Show work for any calculations.

ANALYZE & APPLY QUESTIONS 1. What differences in physical properties allowed you to separate the components of your mixture?

2. Explain how you would separate an alcohol and water mixture.

3. The Chemistry labs use an ion-exchange chromatography resin to make de-ionized water from tap

water. Explain how this method of separation works (you may need to do some research—please

list references if you use them!!!).

CONCLUSION Complete a standard conclusion and error analysis – see “Writing the Lab Report” for help.

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3. COMPARING PHYSICAL AND CHEMICAL CHANGES

INTRODUCTION It is often difficult to tell the difference between physical and chemical changes, but the test of

a physical or chemical change is this: Does the change alter the nature of the substance? Is

something new with brand new chemical and physical properties formed? Were bonds broken or

formed to form a new substance, or is it still the original substance?

If the shape, size, or physical state are altered, but the chemical composition remains the

same, the change is a physical change. In contrast, in a chemical change or reaction, the atoms of a

substance are rearranged with new bonds, and a new substance is formed with a chemical

composition that is different from the chemical composition of the original substance. Chemical

changes can occur simultaneously with physical changes.

In this laboratory activity, you will carry out seven procedures, record observations, and

determine in each chase whether a chemical or physical change (or both!) occurs.

BACKGROUND RESEARCH 1. Explain the difference between qualitative and quantitative data – is this lab primarily quantitative

or qualitative?

2. List five observable indications or signs that a chemical change has occurred.

3. List two examples of an everyday chemical change.

4. List two examples of an everyday physical change.

5. When you dissolve sugar in water to make a sugar-water solution, does the solution still taste like

sugar and water? Based on your answer, do you think making a solution is a chemical or physical

change? Explain.

6. When you bake (or burn, if you’re a terrible cook) cake batter, does the baked/ burnt cake still

taste like the original cake batter? Based on your answer, do you think baking or burning food is a

chemical or physical change? Explain.

SAFETY CONCERNS Wear your goggles and aprons! You may want to wear gloves for this lab. Exercise caution to prevent

burns. Silver nitrate (AgNO3) solution stains skin brownish – you won’t see the stains until later.

Sodium hydroxide (NaOH) solution is corrosive – rinse with water if you get it on your skin. Copper

(Cu) and copper compounds can be irritating to the skin.

MATERIALS Read through the lab procedures and compile a list of reagents and equipment needed.

PROCEDURES For each of the following experiments, make observations and determine if chemical or physical

changes are occurring. There are several steps for each experiment, make sure you determine if each

is chemical or physical.

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#1. (Demonstration). Heating toothpicks over a Bunsen burner.

The instructor will break several toothpicks into small pieces and put the pieces into a large test tube.

The instructor will then heat the tube strongly over a laboratory burner (in the fume hood) until a

change is visible.

CLEANUP: After cooling, the contents will be disposed in a solid waste container.

#2. Combining copper (Cu) wire with silver nitrate (AgNO3) solution.

Mass a piece of copper (Cu) wire, and record the mass. Place the wire in a small test tube and add

silver nitrate (AgNO3) solution until it is just covered. Place the test tube in a rack, and let it stand until

fifteen minutes before the end of the lab since this experiment will take time to obtain any measurable

change. Go on to the next experiment, and return to the next step at the END OF THE LAB. *

*To be done at the END OF THE LAB: Remove the wire with tweezers and scrape and wash off any

material sticking to the wire. Dry the wire with a paper towel. Mass the wire, and record the mass.

CLEANUP: Pour the silver nitrate (AgNO3) solution into the liquid waste container (NOT THE SINK), and

put the copper (Cu) wire in the solid waste container.

#3. Combining sodium chloride (NaCl) and water.

Take a small scoop of sodium chloride (NaCl) with your scoopula, and place it into a small test tube.

Add enough tap water to dissolve the sodium chloride (NaCl), and stir. Keep the sodium chloride

(NaCl) solution for your next two experiments.

#4. Heating the sodium chloride (NaCl) solution over a Bunsen burner.

Set up a ring stand over a Bunsen burner, and mount the iron ring on the ring stand so that the tip of

the outer blue flame would touch the bottom of the ring if the Burner were lit. Making sure that your

Bunsen burner is turned off, place a wire mesh over the iron ring. Place an evaporating dish on the

wire mesh, and pour HALF of the sodium chloride (NaCl) solution from experiment #3 into the

evaporating dish. Keep the other half for your next procedure. Use your Bunsen burner to heat

carefully until all the water is gone—if the contents start to bubble or spatter, remove the burner and

place a watch glass (curved side down) on the evaporating dish. Once covered, continue heating gently

until all the water is gone. Once done, remove the evaporating dish using tongs or an insulated glove

to a ceramic tile.

CLEANUP: Let the dish cool completely, then wash and dry the evaporating dish. If you do not wait

until the dish is cool, it will crack when it comes in contact with the cool tap water!!! Keep the ring

stand setup for experiment #6.

#5. Combining sodium chloride (NaCl) solution with silver nitrate (AgNO3) solution.

Take the test tube containing the other half of the sodium chloride (NaCl) solution from experiment #3.

Add silver nitrate (AgNO3) solution several drops at a time until a change is observed.

CLEANUP: Pour the resulting solution into the liquid waste container (NOT THE SINK).

#6. Heating copper (II) sulfate pentahydrate (CuSO4 � 5 H2O) over a Bunsen burner.

Use the ring stand setup from experiment #4. Take a small scoop of copper (II) sulfate pentahydrate

(CuSO4 � 5 H2O) with your scoopula. If the crystals are very large, use a mortar and pestle to pound/

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grind the crystals into smaller pieces. Place the crystals in an evaporating dish. Heat the solid gently at

first, then more strongly. Once done, remove the evaporating dish using tongs or an insulated glove to

a ceramic tile.

CLEANUP: Let the dish cool completely. Take a very large beaker and fill it with tap water. Put the

contents of the dish into the beaker – be careful! This will generate some heat. Pour all substances

down the sink with plenty of running water, and clean and put away the evaporating dish and beaker.

If the ring stand setup is still hot, please leave it out.

#7. Combining aluminum (Al) with sodium hydroxide (NaOH) solution

Take a small square of aluminum (Al)—about the size of your thumb, and tear it into small pieces (the

smaller the pieces, the faster the next step is). Pour about 100-mL of sodium hydroxide (NaOH)

solution into a 250-mL beaker. Before adding the aluminum, measure and record the temperature of

the solution. Then, add the aluminum (Al), and stir until any change is completed. Continue to

observe the mixture and measure temperature for at least five minutes, recording the highest

temperature reached.

*While you are waiting, go back and finish experiment #2.

CLEANUP: Pour the resulting solution from experiment #7 into the liquid waste container (NOT THE

SINK). If the ring stand setup from experiment #6 is still hot, please leave it out. If is sufficiently cool,

break it down and put it away. Clean all glassware and lab utensils and return them to your locker. If

you are missing equipment or have extra equipment, please let your teacher know.

DATA 1. Calculate the mass change in experiment #2 and the temperature change in experiment #6.

2. Make a data table summarizing the qualitative observations for each experiment.

3. Make a data table summarizing any quantitative data for any of the quantitative experiments.

ANALYZE & APPLY QUESTIONS 1. List three physical properties you observed in the lab. Be specific.

2. Answer the following questions in chart form.

• Which experiments were qualitative and which were quantitative?

• Which changes are physical and which are chemical?

Experiment # qualitative/ quantitative physical/ chemical

CONCLUSION Complete a standard conclusion and error analysis – see “Writing the Lab Report” for help. Make sure

you explain the difference between a physical and chemical change. Use two examples of physical

changes and two examples of chemical changes from this lab to illustrate the difference.

17

4. DETERMINING AN EMPIRICAL FORMULA

INTRODUCTION In a sample of a compound, regardless of the size of the sample, the number of moles of one

element in the sample divided by the number of moles of another element in the sample will form a

small whole-number ratio. These ratios can be used to determine the subscript in the empirical

formula of the compound.

In this experiment, a pure metal—magnesium, will be massed and then reacted with oxygen in

the air to form a binary compound. From the masses before and after the reaction, you will calculate

the moles of each element and find the empirical formula.

BACKGROUND RESEARCH 1. What are the reactants—names and chemical formulas?

2. What product is formed?

3. Will the mass of the contents in the crucible increase or decrease after heating? Explain.

4. How will you determine the mass of oxygen reacted?

5. What would happen to the calculated mass of oxygen reacted in the following cases:

a. If the magnesium is not heated completely?

b. If the magnesium was not properly cleaned and is already partially oxidized?

6. Prepare a data table – include ALL masses that you will need to measure in the course of this lab.

SAFETY CONCERNS Wear your goggles and aprons! Use tongs to handle crucible and cover. Do not place metal directly in

the flame.


PROCEDURES 1. Setup a ring stand with an iron ring and pipe stem triangle at the height of blue inner flame.

2. Clean and dry a crucible and cover. Drive off any remaining moisture by heating the crucible (with

cover ajar) in the hottest part of the burner flame for a few minutes. Allow the crucible and cover

to cool, then mass and record the empty crucible and cover..

3. Clean the length of magnesium ribbon given to you with steel wool until it is shiny, and then cut it

into small pieces. Place the pieces in the crucible, then mass and record the crucible, cover, and

magnesium.

4. Cover the crucible, and place it in a clay triangle. Heat for a few minutes gently—if you see white

smoke escaping, stop heating immediately, and wait until the smoke dissipates. Using the tongs,

carefully tilt the cover and check to see if the magnesium is still shiny (unreacted). If it is no longer

shiny, keep the cover ajar, and heat for about 10 minutes until the contents are completely ashy.

Make sure that no unreacted magnesium remains.

18

5. Use the tongs to CAREFULLY remove the cover and crucible from the ring stand to a tile. Add a few

drops of water slowly, then continue to add just enough water to just cover the contents of the

crucible. Waft to observe any odor – record.

6. Replace the crucible on the ring stand setup, partially covering the crucible with a cover to avoid

splattering. Holding the burner in your hand, gently heat the contents of the crucible by moving

the burner slowly back and forth.

7. When the liquid has boiled off, repeat steps 5 and 6.

8. When all the liquid has boiled off a second time, strongly heat the uncovered crucible for about 5

minutes to make sure the compound is completely dry.

9. Turn off the burner, and use the tongs to CAREFULLY remove the cover and crucible from the ring

stand to a tile. Allow the crucible, cover, and contents to cool, and then mass and record the

crucible, cover, and contents.

DATA & CALCULATIONS 1. Show all calculations needed to determine the empirical formula.

2. Include a % error for the Mg:O ratio.

3. Make a data table for any measurements and calculated numbers.

ANALYZE & APPLY QUESTIONS 1. In addition to the magnesium oxide, a small amount of another magnesium compound forms

during the heating in air. In step 5, water was added to the contents in the crucible to convert this

compound into magnesium oxide. The odor indicated the ammonia was also formed in the

reaction with water. What was the formula of the other magnesium compound? Answer by

writing and balancing the equation:

. ? + water � ammonia + magnesium oxide

2. There are three main sources of error in this lab. How would the ratio of Mg:O be different than

1:1 if (Hint: remember that you want to consider the impact on the calculated mass of oxygen):

a) Some of the product formed was lost in heating at the end (before measuring the mass)?

b) All of the magnesium had not been completely reacted?

c) All of the water had not been completely evaporated?


you give your experimental ratio, and discuss the error in YOUR experiment. (Hint: make sure that the

error is consistent with the results you obtained).

19

5. COMPOSITION OF HYDRATES

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

hydration) as part of their structure. The water is chemically combined with the salt in a definite

ratio. Ratios vary in different hydrates, but are specific for any given hydrate. The formula of a

hydrate is represented in a special manner. The hydrate of cobalt (II) chloride hexahydrate has the

formula

CoCl2 � 6 H2O. The unit formula for the salt appears first, and the waters of hydration follow. The dot

(�) means that the water is loosely bonded to the salt. The coefficient 6 stands for the number of

water molecules bonded to one unit of salt. This formula, like the formulas of other compounds,

illustrated the law of definite composition, and for this particular hydrate, the ratio of salt to water is

1:6.

When hydrates are heated, the chemically bonded “waters of hydration” are released as

vapor. The remaining solid is known as the anhydrous salt or anhydrate. The general reaction is:

heat

Hydrate � Anhydrous salt + Water

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

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

lost by the hydrate. The percentage of water in the original hydrate can easily be calculated.

In this investigation, you will determine the % of water in an unknown hydrate and use it to

identify the hydrate from a given list of hydrates.

BACKGROUND RESEARCH 1. Rewrite the equation in the introduction, including signs for physical state (s,l,g,aq). How will the

mass of the material in the evaporating dish compare before and after heating?

2. State a purpose.

3. Write a procedure. You will be given a sample of unknown hydrate. Describe in detail the

procedure you will use to determine the % water in the hydrate to identify the unknown from the

list below. Write you procedures in list format.

4. List materials needed.

5. Prepare a data table for all measurements.

6. Calculate (and show work) to determine the theoretical %water in each of the possible unknown

hydrates:

a. Cobalt (II) chloride hexahydrate

b. Copper (II) sulfate pentahydrate

c. Strontium chromate monohydrate

d. Calcium chloride dihydrate

e. Barium chloride dehydrate

f. Magnesium sulfate heptahydrate

g. Copper (II) chloride dihydrate

SAFETY CONCERNS Wear your goggles and aprons! Exercise caution to prevent burns.

20

MATERIALS & PROCEDURES See “Background Research” #3-4.

DATA & CALCULATIONS 1. Show all calculations needed to determine the % water in the hydrate.

2. After selecting one of the seven possible hydrates, calculate a % error for the % water.


ANALYZE & APPLY QUESTIONS 1. If all the waters of hydration were not released, how would your calculated % of water compare to

the accepted value? Explain.

2. What could cause the opposite results (from the above question) to occur?

3. When water is added back to the anhydrous salt, is the reaction endothermic or exothermic?

Explain.


you identify your hydrate and support your identification using data. State your percentage error, and

indicate the errors that could have led to your (overly low or high) experimental percentage. Briefly

discuss the principle used to accomplish the lab.

21

6. DETERMINING THE THICKNESS OF ALUMINUM FOIL

INTRODUCTION

This activity mirrors the problem solving done by scientists and engineers. You must work

cooperatively, and all members must be involved in finding a way to accomplish your objectives. You may use

any book or piece of equipment that is present in the room, but you may not search online for answers. You

must also adhere to the rules. You will NOT have a prelab or a regular lab report. Instead, you will perform

and document this lab in one session—you will not be bringing anything home. Some information that may be

helpful. The density of aluminum is 2.69 g/ cm3, and the diameter of one atom is 2.50 x 10-8cm (Assume the

atom is a sphere).

OBJECTIVES: 1. To determine the thickness of a piece of aluminum foil to three significant figures. Use cm as your unit.

2. To determine the thickness of the foil in atoms.

3. To determine the total number of atoms of aluminum in the foil.

4. To answer the questions at the end of this lab.

SOME DIRECTIONS: 1. Write the names of the group members on the top of a piece of looseleaf.

2. Read the following RULES, and adhere to them during the lab!!!! Violations will impact your grade!

• You may NOT destroy, crumble, tear, or fold your foil. It must be returned in exactly the same condition that you

received it.

• You can only confer with members of your group. You are NOT permitted to spy or eavesdrop on other groups—I

should not even see you near another table or other group members!

• You must turn in all materials at the end of the period—you should not remove any papers from this lab!

3. On another piece of paper, do some brainstorming. Maybe draw a sketch of your aluminum foil, or write

down all the information or formulas that may be related to the objectives. Remember – you may need to

bring in previously learned concepts to complete this lab.

4. Make sure everything is cleaned up at the end of your lab!

WHAT IS TO BE SUBMITTED BY EACH GROUP AT THE END OF THE PERIOD:

The aluminum foil and one handwritten report with: a. (5) Group member names

b. (5) Paper with brainstorming

c. (10) Procedure (list format)

d. (10) Data table

e. (20) Lab calculations with answers boxed

f. (10) Answers to questions with answers

boxed

g. 10 points will be assessed for neatness.

QUESTIONS: 1. If the population of the world is about 5.6 x 109 people, how many atoms of aluminum from your foil

could you distribute to each person in the world?

2. If the cost of aluminum foil is $5.49 per 75 square fet, calculate the cost of your piece of foil.

3. If aluminum foil with the same thickness as your piece is made of an alloy that is only 98.5% aluminum

by mass, and the roll of foil has an area of 75 square feet, calculate the number of aluminum atoms in

the roll.

4. If you had a piece of gold foil with the same dimensions as your piece of aluminum foil, what would be

its mass in grams? (Density of gold = 19.3 g/ cm3)

5. If the atomic radius of gold is 144 pm, calculate the number of gold atoms in your place.

GRADING:

Performance/ adherence to rules: 30 pts Writeup: 70 pts (10 pts – neatness/ formatting)

22

7. TYPES OF CHEMICAL REACTIONS

INTRODUCTION Have you ever thought about the fact that aluminum, unlike iron, doesn’t seem to rust?

Actually, aluminum does react with the oxygen in the air to form the dull, white compound aluminum

oxide, Al2O3 (reaction 1). The aluminum oxide remains on the aluminum as a protective coating, and

no further reaction occurs.

Have you ever wondered why hydrogen peroxide (H2O2) is sold in brown bottles? Hydrogen

peroxide decomposes gradually into water and oxygen, but light hastens the reaction (reaction 2).

In 1776, Henry Cavendish reacted hydrochloric acid with zinc to produce hydrogen gas. Zinc

chloride remained behind in the vessel he used (reaction 3).

Hydrochloric acid reacts with sodium hydroxide to produce table salt and water (reaction 4).

Once you have learned to write a balanced equation with it correct reactants, products,

formulas, and coefficients, you have crossed one of the major hurdles or chemistry.

In this experiment, you will observe nine different reactions, identify each type of reaction,

and write a balanced equation for it.

Note that you may use a glowing or lit splint to test for the presence of hydrogen (the splint

ignites the hydrogen with a popping sound), oxygen (the splint’s flame grows), and carbon dioxide (a

lit splint is extinguished). This test must be done immediately, or the gas escapes, and the test will

not work.

Please do not try to dry the inside of test tubes with paper towel – just use a new, dry test

tube.

BACKGROUND RESEARCH 1. What are the five major types of reactions, and what is the distinguishing characteristic for each?

2. Read the introduction. Write balanced equations for the four reactions described in the

introduction above, and identify the reaction type.

3. Read the procedures, write the reactants (chemical formulas) for each of the nine reactions. If you

can predict the products and the type of reaction, feel free to also do that, but you will obviously

have to perform the experiment to make your observations.

SAFETY CONCERNS Wear your goggles and aprons! Use test tube holders or insulated gloves to handle hot test tubes, and

remember to never point open test tubes toward yourself or others. Hydrochloric acid is corrosive –

protect yourself during the lab and during cleanup.


PROCEDURES As you carry out the following reactions, record all your observations.

23

REACTION A.

1. Use your scoopula to place copper (II) carbonate in a small test tube to a height of about 1 cm.

Note the appearance of the sample.

2. Wearing an insulated glove and using the test tube holder (not clamp), heat the tube strongly until

a change occurs. Quickly remove from flame, and immediately hold a burning splint in the mouth

of the test tube to test for gas. Observe the effect on the splint and the appearance of the sample

in the test tube.

3. CLEANUP: Discard splint and contents of tube in a paper towel and dispose into the waste basket.

Once test tube has cooled, use water and a test tube brush to scrub the test tube clean. Rinse and

dry scoopula clean.

REACTION B.

1. Repeat the same procedures done for reaction A (using a new dry test tube), but with ammonium

carbonate.

2. Waft the test tube, and describe/ record any readily apparent odor.

3. Hold a piece of red litmus paper to the mouth of the test tube, and observe/ record.

4. CLEANUP: Discard splint and litmus paper into the waste basket. Once test tube has cooled, flush

contents of tube down sink with water. Use water and a test tube brush to scrub the test tube

clean. Rinse and dry scoopula clean.

REACTION C.

1. Place a very small amount of calcium carbonate in a dry test tube.

2. Add about 20 drops of hydrochloric acid to the test tube. As soon as the reaction is observed, hold

a burning splint in the mouth of the test tube.

3. CLEANUP: Discard splint in the waste basket. Flush contents of tube down sink with water. Use

water and a test tube brush to scrub the test tube clean. Rinse and dry scoopula clean.

REACTION D.

1. Obtain a piece of magnesium ribbon, tongs, and a watch glass.

2. Using the tongs, place the ribbon in the flame. Remove it as soon as it begins to burn, and hold it

over the watch glass. DO NOT LOOK DIRECTLY AT THE BURNING MAGNESIUM.

3. Examine the residue.

4. CLEANUP: Rinse tongs and watch glass clean, and dry.

REACTION E.

1. Place about 1 mL of sodium chloride solution into a small test tube.

2. Add silver nitrate solution to the sodium chloride a few drops at a time until you see a visible

reaction.

3. CLEANUP: Pour contents of test tube into the designated liquid waste container, and then use

water and a test tube brush to scrub the test tube clean.

REACTION F.

24

1. Add about 20 drops of sodium hydroxide solution into a small test tube. Add one drop of

phenolphthalein (an acid-base indicator) and mix by gently swirling the tube.

2. Add hydrochloric acid one drop at a time to the test tube. Count and record the minimum number

of drops of acid required for a permanent color change.

3. CLEANUP: Add a few drops of sodium hydroxide to the test tube until contents are pale pink (NOT

magenta). Flush contents of test tube down sink with water. Use water and a test tube brush to

scrub the test tube clean.

REACTION G (see figure to right).

1. Place about 10 mL of hydrochloric acid in a

large test tube. Insert the one-hole stopper

with the glass tube/ rubber hose into the test

tube, and use a test tube clamp to clamp

onto a ring stand. CAUTION: Hydrochloric

acid is corrosive!

2. Fill a small test tube with tap water, and

invert the tube in a trough or large beaker

filled with enough tap water so that your

hand can fit under the test tube. Stand the

inverted test tube in the trough – you will put

the hose into the test tube to collect gas

later.

3. Obtain one piece of mossy zinc, and open the stopper to carefully drop it into the hydrochloric acid.

Immediately cover the tube with the stopper. Let the gas generated displace the air in the reaction

test tube and hose for about 30 sec to a 1 minute before using your hand to put the hose into the

gas-collecting test tube. Observe the reaction occurring in the medium test tube.

4. Allow the collection tube to fill with gas. Collect another if gas is still being produced. Stopper the

tubes of gas with solid stoppers and remove from the water, keeping them inverted.

5. CLEANUP: You will clean up this reaction after REACTION H.

REACTION H.

1. Keeping the gas-filled test tube inverted, dry off the outside, and then remove the stopper.

2. Carefully bring a burning splint to the mouth of the gas-filled tube. Notice any change in the tube,

and record your observations.

3. CLEANUP: Carefully dispose of the acid from REACTION G in the designated liquid waste container.

BE SURE TO KEEP GOGGLES ON DURING CLEANUP! Use water and a test tube brush to clean the

test tubes.

REACTION I (demonstration).

1. Your instructor will take some ethanol and ignite it. Observe the reaction.

DATA & CALCULATIONS 1. Make a data table summarizing your observations, the reaction type, and a BALANCED chemical

equation for the nine reactions.

25

ANALYZE & APPLY QUESTIONS 1. What gases were generated in this lab? How did you identify them? Be specific.

2. Name a characteristic that all the gases have in common. What characteristic distinguishes them

from each other?

3. Identify the reactions that were redox (reduction/ oxidation).

4. What was The Hindenburg, and what is significant about it? Which reaction in this lab is related the

even concerning the Hindenburg? Write the balanced chemical equation.


26

8. RELATING MOLES TO COEFFICIENTS OF A CHEMICAL EQUATION

INTRODUCTION Coefficients in a chemical equation indicate the number of moles of each substance;

therefore, the ratio of moles of a substance to moles of any other substance in the reaction can be

determined at a glance.

In this experiment, iron filings will be added to an aqueous solution of copper (II) sulfate and a

reaction will take place. The copper produced will be separated from the solution, dried, an dmassed.

From the known masses of reactants and products, the moles of each can be calculated and the

equation written.

You will compare the experimental mole ratio to the coefficient ratio of the possible balanced

equations. After selecting one equation, you will also use it to calculate the theoretical yield and

determine the percentage yield of copper from your experimental data.

BACKGROUND RESEARCH 1. Calculate how many grams of copper (II) sulfate pentahydrate must be massed to give 8 grams of

anhydrous copper (II) sulfate. This reactant will be in excess to ensure that all the iron reacts.

2. What type of reaction occurs when iron and copper (II) sulfate are reacted? Write the two possible

complete balanced chemical equation (one with Fe+2 and one with Fe+3).

SAFETY CONCERNS Wear your goggles and aprons!


PROCEDURES 1. Obtain a 100-mL beaker, and clean it thoroughly to remove any material clinging to the walls. (The

beakers are used only for this experiment, and there may be a stain that is not removable). Dry the

beaker, label it with your group’s name, and RECORD THE MASS of your empty beaker.

2. Measure the proper amount of copper sulfate pentahydrate (amount you calculated) and add it to

the beaker. Record the exact mass added.

3. Measure about 50 mL distilled water and add to the beaker.

4. Mass about 2.24 g of iron filing on a piece of paper. Record the exact mass and set aside.

5. Heat the beaker using a burner/ ring stand setup or a hot plate to just below boiling, stirring until

the crystals completely dissolve. DO NOT ALLOW THE LIQUID TO BOIL – you are simply heating the

liquid to increase the speed at which the crystals dissolve and the reaction proceeds.

6. Using insulated gloves, remove the hot beaker and place on a tile. SLOWLY, A LITTLE AT A TIME,

add the iron filings to the hot solutions. You want to avoid adding the iron all at once because this

will cause a sudden burst of vapor. After all the iron has been added and the mixture stirred, allow

the beaker to sit for at least 10 minutes or until the reaction is complete.

7. Decant the liquid into the waste beaker using your stirring rod to avoid splatters (as demonstrated).

Try not to disturb the solid (i.e., copper) at the bottom of the beaker.

27

8. Add about 10 mL of water to the solid remaining in the beaker. Stir vigorously to wash the solid

and get rid of any unreacted copper (II) sulfate. Let it settle, and decant the liquid into the waste

beaker.

9. Repeat step #8 at least 6 times to ensure that only copper and some water remain in the beaker. If

there is any blue/ green residue visible when your solid dries overnight, you did not wash your

copper adequately.

10. Spread the solid out on the bottom of the beaker and place it on the tray provided to dry over

night.

11. When it is dry, find the mass of the beaker and copper metal.

12. CLEANUP: Scrape out contents in the waste basket, RETURN THE BEAKER as directed.

DATA & CALCULATIONS 1. Calculate the moles of iron reacted and copper formed.

2. Write the two possible balanced equations, and calculate the two possible Fe: Cu coefficient ratios.

3. Compare your experimental Fe: Cu molar ratio with the two possible coefficient ratios, and select

one equation. Use this equation to calculate the theoretical yield of copper using the reacted iron

as the limiting reactant.

4. Calculate the percentage yield of copper.


ANALYZE & APPLY QUESTIONS 1. Would this reaction occur with copper metal and iron sulfate as the reactants? Explain.

2. How do you know whether iron (II) or iron (III) sulfate forms? Use your data and calculations to

explain.

3. What impact do the following errors have upon your experimental Fe: Cu molar ratio? Explain

briefly.

a) You spilled some of the iron filings out while pouring the iron into the hot copper (II) sulfate

solution.

b) You did not properly wash the Cu fully, and there is still some unreacted copper (II) sulfate

mixed in.

c) The Cu was not fully dried out before you did your final mass measurement.


you explain how your experimental Fe: Cu molar ratio compares to the coefficient ratio of the balanced

equation you selected. Refer to your calculations. Also, discuss the factors that caused YOUR yield to

be either too high or too low.

28

9. PARTICLE SIZE FROM COLLISION PROBABILITIES

INTRODUCTION There are situations where it is necessary to use indirect means to gather information or make

measurements. For example, scientists frequently use indirect means to study galaxies, the size of

the universe, and the atom.

In this lab activity, you will use probability to determine the size of a marble. By counting how

often you succeed in hitting a marble when another marble is shot at a group of them, you can apply

a formula that will calculate the probable size.

After getting an indirect measurement of your marble, you will then get a direct measurement

of the actual size of the marble, and compare your indirect measurement with the actual size.

BACKGROUND RESEARCH 1. Sketch and describe Rutherford’s experiment. What did he conclude about atomic structure?

2. How is this experiment a model of Rutherford’s experiment?

3. Summarize the rules regarding marble placement.

4. A meter stick is marked every 0.1cm, to what decimal place should measurements using it be

made?

SAFETY CONCERNS You do not need to wear goggles or gloves, but please make sure that you are aware of rolling marbles

as you walk around.


PROCEDURES 1. First, you will imagine that you cannot directly measure the size of a marble, so you will be

determining its size indirectly (through probability). Obtain ten target marbles and one additional

marble for bombarding the target marbles. Using sticky tack to keep the marbles where you have

place them, arrange the ten target marbles inside a three sided enclosure made from three meter

sticks following the following rules:

a. Try to place target marbles randomly.

b. There must always be room for the bombarding marble to get through (i.e., don’t place the

marbles too close to each other!).

c. Do not place marbles in front of each other so that they are “shielding” other marbles from

being hit.

2. Randomly roll (not throw) the bombarding marble toward the target marbles along a path

perpendicular to the target marbles and parallel to the side walls of the enclosure. Do not aim. It

may help to close your eyes. Roll a minimum of 300 times, counting and recording both the total

number of rolls and the number of hits. A “hit” occurs when the bombarding marble hits another

marble before hitting the wall – secondary hits of other marbles do not count. RESULTS IMPROVE

AS THE NUMBER OF ROLLS INCREASES.

29

3. Now that you have done an indirect measurement, obtain a direct measurement by lining up your

ten target marbles (NOT your extra bombarding marble) between two meter sticks in a straight line

so that each marble touches another. Record the length of ten marbles.

4. CLEANUP: Return 11 marbles in a beaker with all of the sticky tack stuck to the inside of the beaker.

Return the meter sticks.

DATA & CALCULATIONS 1. Indirectly calculate the diameter (length) of a marble to 3 significant figures from the experimental

data and a probability based formula:

D = H x d where D = marble diameter

2N x Tr H = number of hits

d = length of enclosure = 100.00 cm

N = number of target marbles

Tr= number of total rolls (trials)

2. Using your measurement of the length of ten marbles, determine the diameter (length) of one

marble (in cm).

3. Calculate the percent error for your indirect measurement of marble diameter. Assume the marble

diameter obtained from the direct measurement as your actual value.


ANALYZE & APPLY QUESTIONS 1. Explain how the following errors would impact (increase or decrease) your calculated marble

diameter:

a) Aiming to hit marbles.

b) Placing two marbles so that one blocks another from being hit.

c) Placing two marbles so that they are touching.

2. Besides using larger, easy-to-visualize materials, what is the most significant way in which this

model of Rutherford’s experiment differ from the actual experiment?


30

10. FLAME TESTS AND EMISSION SPECTROSCOPY

INTRODUCTION When an element’s gas or ion form is energized with heat or electricity, it gives off

electromagnetic radiation or light. If the light emitted by a gas is passed through a spectroscope, a

pattern of narrow bands of light is produced. Each element produces its own distinct pattern that

differs from the pattern of every other element, which is referred to as that element’s emission

spectrum or bright-line spectrum. Because the emission spectrum for each element is unique, it can

be used for element identification, just as fingerprints (or DNA) can be used to identify an individual.

When an electron absorbs energy, it moves to a higher “excited” energy level from its lowest

allowed energy level (i.e., “ground state”). This “excited” electron cannot remain at any of these

higher allowed energy levels if there are unoccupied lower energy levels closer to the nucleus, and as

the electron relaxes to a lower energy state, a photon is given off that has an energy equal to the

energy difference between the higher and lower energy levels. Each band in the bright-line spectrum

corresponds to a specific electron transition, and the energy of these transitions (and of the emitted

photons) is given by Max Planck’s equation, E = hv, where E = energy in Joules, v = frequency, and h =

Planck’s constant, 6.63 x 10-34 Joule-seconds.

In Part 1, you will energize a small amount of several metallic salts (the ion form of each

metal). Each element will give off a unique set of photons with specific wavelengths and colors, but

without a spectroscope, we do not see each color emitted. Instead, your brain recognizes several

photons with different wavelengths as one particular color. You will perform a flame test on several

metallic salts, and record the flame color for each metal ion (being as specific as possible). You will

then perform a flame test on an unknown metallic salt and identify it. You will use cobalt glass as a

tool for reducing the “masking” effect of one element upon another element in a mixture.

In Part II, you will use a spectroscope to examine the bright-line spectra of some gaseous

elements and the continuous spectra from other light sources. You will also identify an unknown by

comparing its observed spectrum to the spectrum chart.

BACKGROUND RESEARCH 1. What is meant by the term “ground state”?

2. How do electrons become “excited”?

3. When is the energy absorbed by electrons released?

4. What is the form of energy released?

5. State the equation that is used to determine the energy content of a packet of light of specific

frequency.

6. How should the burner flame be adjusted for best results?

SAFETY CONCERNS Wear goggles and aprons, and be aware of loose clothing and long hair (it should be tied back!) around

the Bunsen burners. Do not touch the power supplies or tubes on the discharge tubes.


31

PROCEDURES

PART 1 1. Obtain a set of wooden splints (take only as much as you need), and soak them in DISTILLED water.

Tap water contains dissolved metallic salts, which would contaminate your samples and distort

your results. Obtain your metallic salts (7 plus 1 mixture): BaCl2, CaCl2, CuCl2, LiCl, KCl, NaCl, SrCl2,

and a NaCl/ KCl mixture.

2. Set up a separate beaker of tap water for extinguishing your burning splints.

3. Adjust the Bunsen burner to a blue flame.

4. Take your wet wooden splint. Dip your splint into either a solution of your metallic salt or into a

solid sample of your metallic salt – you only need a small crystal or some solution on the end of

your stick to get your results. If you take too much, you will contaminate the burner with residual

salt.

5. Hold the splint at the top of the flame, and record the color – do not mistake the color emitted

with the color of a burning splint. Record the color as specifically as you can, and repeat with all of

the different metallic salts and the sodium/ potassium mixture.

6. Since the sodium and potassium flames can look similar, record how the flame looks when seen

through a cobalt glass. Then see how the flame of the Na/ K mixture looks with the cobalt glass.

7. CLEANUP: Discard all used splints into the garbage can. Dump any water or waste solution in the

beakers down the drain, and rinse the beakers clean. Return beakers and Bunsen burners to your

lockers, and make sure you lab bench is wiped clean (use a sponge).

PART II

1. Using the spectroscopes, view the 5 discharge tubes located in the room. Record what you see--

sketching the location and color of the bands of light observed (use colored pencils).

2. Compare the observed bands to the known spectra of several elements, and identify the neon

tube.

3. CLEANUP: Return spectroscopes and colored pencils. There should be no stray pencils.

DATA & CALCULATIONS 1. Make a data table summarizing your observations of flame color from Part 1 – don’t forget to

include observations made with the cobalt glass.

2. Include a data table or figure showing your observations from Part 2.

ANALYZE & APPLY QUESTIONS 1. Give two reasons why a flame test might not always give you a valid element identification.

2. Explain why potassium was visible in the mixture when seen through the cobalt glass.

3. Can you identify EACH element in a mixture, with a flame test? From an emission spectrum?

4. What is the difference between an absorption and an emission spectrum? Which is observed from

the discharge tubes?

5. What would be observed if a spectroscope were used during a flame test?

32

CONCLUSION Complete a standard conclusion and error analysis – see “Writing the Lab Report” for help. Don’t

forget to identify your unknown in Part 1, and identify which tube contained the element to be

identified. Support your conclusion with your data.

11. INVESTIGATION OF THE HYDROGEN SPECTUM

INTRODUCTION The experimental emission spectrum of hydrogen has discrete bands of particular energy.

Bohr knew that the correct atomic model had to account for the experimental spectrum of

hydrogen. His model assumed that the single electron in the hydrogen atom could be located in a

particular orbit and would remain there unless excited by energy from outside the atom. Each orbit

has a particular energy associated with it, given by the equation:

En = - RH (1/n2) where n = principal quantum number (integer 1 . . . 7)

RH = Rydberg constant, 2.18 x 10-18

The negative value for the energy is a convention to ensure that the energy of the electron in the

atom is lower than the energy of a free electron, which is assigned a value of zero. As the electron

gets closer to the nucleus (as n decreases), En becomes more negative. When n = 1, this is the

ground state for the hydrogen atom.

Bohr’s restriction on the energies of the orbits meant that the energies associated with

electron motion between permitted orbits would also be fixed in value (“quantized”). The emission

of radiation by an excited hydrogen atom could then be explained in terms of the electron dropping

from a higher energy orbit to a lower one and giving up a quantum of energy (a photon) in the form

of light. As the electron moves, or transitions, between energy levels, the energy absorbed or given

off is equal to the difference in energy levels:

∆E = Einitial - Efinal where Einitial = initial energy level in a transition

Efinal = final energy level in a transition

In this activity, you will:

• Calculate the energies of the possible energy levels (orbits) of the hydrogen atom shown on

the diagram.

• Measure the wavelengths of the lines on the hydrogen spectrum

• Calculate the energy of the photon for each line.

• Determine which energy level transition was responsible for each of the four visible lines.

BACKGROUND RESEARCH 1. Using the equation given, calculate the energies for each energy level (n = 1 to n = 7), and write it

and draw an energy level (properly scaled) on the appropriate place on an energy level diagram.

Make sure n = ∞, which corresponds to a free electron, is labeled, and is the highest possible

energy level. Hint: You may want to include (x 10-20 J) in the units of the y-axis so that you do not

33

need to graph very small decimals (e.g., for the number, 1.23 x 10-19, you could graph 12.3 (x 10-

20)).

2. Look up the hydrogen spectrum, and record the wavelength of the red, blue, and two violet lines in

nanometers and in meters.

3. Calculate the energy (in Joules) of the four lines using Planck’s equation, showing work.

SAFETY CONCERNS There are no safety concerns associated with this lab.

MATERIALS All you need is paper, pen/ pencil, and a calculator (and possibly, a ruler).

PROCEDURES 1. Using the electron energies for each energy level (orbit), determine which energy level transitions

were responsible for each of the four spectral lines on hydrogen’s emission spectrum. Note: There

are many more possible transitions for the electron than the ones you can observe in the visible

spectrum. Show calculations to back up your answers.

DATA & CALCULATIONS 1. Include a figure showing a labeled energy level diagram (for n =1 to n = 7)

2. Include a data table summarizing your wavelengths and calculated numbers.

3. Include a data table summarizing which transitions are responsible for each spectral line (e.g. RED,

from n = ____ to n = ____).

4. Show all calculations.

ANALYZE & APPLY QUESTIONS 1. The possible electronic transitions are

grouped into series and fall into different

regions of the electromagnetic spectrum,

depending on the amount of energy

released. The diagram to the right shows

all the possible transitions, and the names

of the series. Determine which series you

observed.

CONCLUSION Complete a standard conclusion and error

analysis – see “Writing the Lab Report” for

help. Support your conclusions with data.

34

12. DETERMINING THE HALF-LIFE OF Ba-137m

(A class experiment)

INTRODUCTION The half-life of a radioactive isotope (radioisotope) is the time required for the activity (or

mass) of the sample to be reduced to one half of its original activity (or mass). In this experiment,

you will determine the half-life of the radioisotope Bm-137m. The m stands for “metastable,” a

condition of temporary stability. A Ba-137 generator will be used as the source of the Ba-137m. The

generator contains Cs-137, which has a half-life of over 30 years and decays by beta emission to Ba-

137m. The Ba-137m that is produced possesses more energy than is normally possessed by a stable

nucleus and “cools” (decays) to stable Ba-137 by gamma emission with a half-life that you will

determine.

BACKGROUND RESEARCH 1. If you were to take a background reading over 5 minutes, how would you calculate the background

counts per minute (cpm)?

2. If your activity readings are taken over 30 seconds, how would you calculate the cpm (gross – not

factoring in the background)?

3. How do you calculate the cpm (net) from the cpm (gross)?

SAFETY CONCERNS There are no serious safety concerns associated with this lab, and the instructor will handle all

materials, making sure to wash hands after the lab has concluded.

MATERIALS Cs-137/ Ba-137m minigenerator Geiger tube and sample holder Geiger counter

PROCEDURES 1. The Geiger counter will be turned on by your teacher and allowed to run without a sample for 5

minutes so that a background radiation count can be taken. RECORD the counts.

2. Your teacher will “milk” the minigenerator to wash out a 7 or 8 drop sample of Ba-137m into the

small metal dish. A Geiger tube will be placed over the sample in a sample holder.

3. The counter will be turned on for 30 seconds and then stopped. You will quickly RECORD the

counts.

4. The counter will then be set to zero and restarted 30 seconds after it was turned off. In this

manner you will collect several 30 second readings with a 30 second “rest” between each reading.

This will continue for about 10 minutes. RECORD your data as the experiment proceeds.

5. CLEANUP: The sample will be flushed down the drain after 10-15 minutes.

DATA & CALCULATIONS 1. Include a data table of your time interval, counts per 30 second interval, gross cpm, and net cpm.

Note the background radiation over 5 minutes and the background cpm.

35

2. Construct a graph of the data. Your cpm (net) should be plotted on the y-axis and the time (either

in seconds or minutes) on the x-axis. Use the beginning of each interval as your time. (For example:

for the interval 0-30 sec, plot your activity versus 0 seconds.) Remember to draw the best fit

smooth curve.

3. Use your graph to calculate the half-life of Ba-137m. Choose any convenient cpm (net) for your

first activity and find the time that corresponds to that activity. Then take half of that first activity

(this is your second activity) and find the time that corresponds to that. Repeat using half of the

second activity. Be sure to do at least two trials and average the results.

4. If the “accepted” value for the half-life is 2.6 minutes, what is your percent error? Show your

calculation. If you plotted your graph in seconds, remember to convert your units

ANALYZE & APPLY QUESTIONS 1. Based on your graph, what is the half-life of Ba-137m? How does it compare to the actual value?

2. Some Cs-137 is always mixed into the Ba-137m sample. Does this affect the measurement of the

half-life of the Ba-137m sample? Explain.

3. Why is it safe to throw the remaining Ba-137 down the drain after 10 or 15 minutes?

CONCLUSION Complete a standard conclusion and error analysis – see “Writing the Lab Report” for help. Support

your conclusion with your data.

36

13. PERIODIC TABLE: A STUDY OF THE REACTIVITY OF METALS

INTRODUCTION The relative reactivity of an element can be predicted by its position on the periodic table. In

this lab, you will determine the trends in reactivity within a group and across a period. In Part 1, you

will observe reactions of several metals in group 1, 2, and 13 with water. In Part 2, you will observe

whether group 2 metals form precipitates. The more active a group 2 metal is, the more precipitate

it will form.

BACKGROUND RESEARCH 1. Write the general single replacement reaction for an active metal M with water, predicting the

proper products.

2. How does phenolphthalein indicate the presence of a base?

SAFETY CONCERNS PART 1: Wear goggles and aprons, and be aware of loose clothing and long hair (it should be tied

back!) around the Bunsen burners (or hotplates). Remember to point the tube away from yourself and

others while heating. You will also be working with calcium – do not touch the calcium with your

hands. It is corrosive to skin!

PART 2: Wear goggles and aprons. Salts of barium and strontium are extremely toxic. Avoid contact

with these chemicals, and wash your hands thoroughly after use.


PROCEDURES

PART 1: REACTIVITY WITH WATER The reactivity of alkali metals (lithium, sodium, and potassium) with water will be observed by

demonstration.

Please note that given time constraints and your groups ability to cooperate, you may decide to test

calcium, magnesium, and aluminum simultaneously.

1. Pour about 5 mL of distilled water into a clean, dry test tube (the “reaction” tube) and place it in

the test tube rack. Using the forceps, add a piece of calcium to the water, and quickly invert a

larger dry test tube (the “gas collection” tube) over the reaction tube to collect any gas being

released. Observe the reaction.

2. Keeping the gas collection tube inverted, test for hydrogen gas with a burning splint inserted in the

mouth of the tube. Record your observations.

3. Add a drop of phenolphthalein solution to the reaction tube. Record your observations

(phenolphthalein indicates the presence of a base.)

37

4. Repeat steps 1 and 2 with magnesium, breaking the magnesium ribbon into smaller pieces before

inserting it into the distilled water. Record observations.

5. If there is no visible reaction for magnesium, remove the gas collection tube, and gently heat the

reaction tube until the water is just boiling, turning off the burner once that point has been

reached.

6. Quickly return the reaction tube to the rack, and cover with the gas collection tube. Carefully

observe if a gas is being formed by looking on the surface of the metal pieces for the presence of

small bubbles. After waiting for a few minutes, test for hydrogen gas again. Record observations.

7. Repeat step 3 with the magnesium in the reaction tube. Remember to record your observations.

8. Repeat the same steps that you performed for magnesium (steps 4-7) with one aluminum pellet.

9. CLEANUP: Remove the aluminum pellet and any magnesium pieces from the water using forceps,

and place in the designated solid waste. Pour any solutions down the sink with plenty of running

water. Clean all test tubes with running water and a test tube brush.

PART 2: PRECIPITATION REACTIONS 1. Obtain dropper bottles containing Mg(NO3)2, Ca(NO3)2, Sr(NO3)2, Ba(NO3)2 solutions.

2. Obtain a 12-well spot plate, and fill 3 wells with 10 drops each for each alkaline earth nitrate

solution (3 wells of Mg(NO3)2; 3 wells of Ca(NO3)2, 3 wells of Sr(NO3)2, and 3 wells of Ba(NO3)2).

3. Add 3 drops of sulfuric acid, H2SO4, to the first well of each nitrate solution. Record whether a

precipitate appears (any cloudiness or film could be a precipitate), the relative amount, and color.

4. Repeat with sodium carbonate, Na2CO3, adding 3 drops to the second well of each nitrate solution.

5. Repeat with sodium chromate, Na2CrO4, adding 3 drops to the third well of each nitrate solution.

6. Test and identify the unknown (one of the metal nitrates you already tested) by reacting 10 drops

of your unknown with 3 drops of H2SO4, Na2CO3, and Na2CrO4, respectively.

7. CLEANUP: Rinse off the spot plate in the sink, using a paper towel or sponge to get rid of any

residue.

DATA & CALCULATIONS 1. Include data tables summarizing your observations from Part 1 and from Part 2.

2. Include a data table with balanced equations for the reactions of the metals and water in Part 1.

3. Include a data table with balanced equations for the precipitation reactions that occurred in Part 2,

making sure to identify the precipitate.

ANALYZE & APPLY QUESTIONS 1. What trend in reactivity with water was observed within metals in Group 1? In Group II? Among

the three metals tested in period 3? (Use < or > in your answer)

2. In comparison to the elements you tested, predict the relative reactivity of the following elements:

Be, Sr, Fr, Zn, and Fe. (Include the elements you tested, and use < or > in your answer and bold Be,

Sr, Fr, Zn, and Fe).

3. Describe any relationships that you can determine between the number of precipitates formed by

each compound and the location of the alkaline earth metal on the periodic table.

38

CONCLUSION Complete a standard conclusion and error analysis – see “Writing the Lab Report” for help. Identify

your unknown in Part 2, backing up your conclusion with your data.

14. PREPARATION AND PROPERTIES OF OXYGEN

INTRODUCTION Oxygen is the most abundant element in nature. While oxygen is 21% of the atmosphere, it

is not readily isolated from the atmosphere under laboratory conditions. It can, however, be readily

generated in the laboratory from oxygen containing compounds (i.e., oxidizers). There are several

methods which are generally used to prepare oxygen: the pyrolysis of a chlorate, the electrolysis of

water, the pyrolysis of a heavy metal binary oxide such as mercuric oxide, and the decomposition of

hydrogen peroxide using enzymes. Pyrolysis is the destructive decomposition of a compound by

heat. All methods are useful in that they are quantitative.

The method to be used in this experiment is the pyrolysis of potassium chlorate. A catalyst,

MnO2, will be used.

This experiment will involve both quantitative and qualitative data. You will prepare oxygen,

collect it by water displacement, and determine the % yield and the % oxygen in potassium chlorate.

You will then study some physical and chemical properties of the oxygen you have collected. You

will write equations for all the reactions.

BACKGROUND RESEARCH 1. What is the purpose of a catalyst? What will be used to catalyze this reaction?

2. How do you collect a gas by water displacement? Sketch the apparatus that you think would be

appropriate for use in this lab.

3. How will you determine the total volume of oxygen collected?

4. What does litmus paper indicate?

SAFETY CONCERNS Wear googles and aprons, and follow all directions. Potassium chlorate is an explosive oxidizer.


PROCEDURES NOTE: You will work in pairs to prepare oxygen. TO complete the reactions in Part 2, you will need a

minimum of 4 bottles of oxygen. Some of you will not collect this much, therefore it may be necessary

to combine groups for Part 2.

PART 1: THE PREPARATION AND COLLECTION OF OXYGEN GAS 1. Record the room temperature and the atmospheric pressure.

2. Mass a clean, dry large test tube.

39

3. Add approximately 3 g of KClO3. TRY NOT TO GET ANY ON THE SIDES OF THE TEST TUBE. WIPE ANY

FROM AROUND THE TOP. Record the mass of the test tube and solid.

4. Add approximately 1.5 g of MnO2. Break up any lumps and mix the contents thoroughly, then

record the mass of the test tube and the contents again.

5. Connect the test tube to an apparatus for the collection of a gas by water displacement. HAVE

YOUR SET-UP APPROVED BEFORE YOU BEGIN THE NEXT STEP.

6. Heat the test tube gently holding the burner in your hand. CONTROL THE HEATING, and do NOT

allow the chemicals to rise in the test tube. If the reaction appears to be going too fast, decrease

the heating.

7. Do not collect the first bubbles that appear, since this will contain mostly air. Collect as many

bottle as you can, covering the bottles with glass plates (oxygen gas is more dense than air so you

can keep the bottles upright).

8. When the bubbling stops, remove heat and IMMEDIATELY DISCONNECT THE TEST TUBE to prevent

water from backing up into the apparatus.

9. Place the test tube in a safe place out of the way to be massed when cool. Proceed with Part 2.

PART 2: PROPERTIES OF OXYGEN Record observations about the physical properties. Record all observations of the following

reactions.

REACTION 1.

1. Place a small amount of charcoal into a deflagrating spoons. Heat until

red hot and quickly lower into a bottle of oxygen. DO NOT ALLOW THE

SPOON TO TOUCH THE SIDES OR BOTTOM OF THE GLASS. Keep the

glass covered as much as possible. See figure to the right. When the

reaction stops, remove the spoon and discard the ash in the waste

pail, keeping the jar covered. If there isn’t any water in the bottom,

add about 10 mL of water and shake. Then, test the water with litmus

paper.

REACTION 2.

2. Heat a small piece of magnesium in a deflagrating spoon until it begins

to glow. Lower it into another bottle following the directions above.

Do not look directly at the jar. When the reaction stops, remove the

spoon. Allow it to cool, and then add a small amount of water to the

SPOON to dissolve the material. Test the solution in the spoon with

litmus.

REACTION 3.

3. Take a small amount of steel wool in the spoon. Heat until glowing,

and again lower into the bottle as in step 1. (No other tests need to be done with the steel wool).

When the reaction stops, discard the contents in the waste pail.

REACTION 4.

4. Place a glowing splint into the last bottle. If it lights up, quickly remove the splint and blow out

flame, covering the bottle. Repeat, and see how many re-lights you can get. Record

40

5. Once done with all reactions, return to mass the cool test tube. Determine the total volume of

oxygen collected in your bottles.

6. CLEANUP: Quench and throw out any splints. Put any used litmus paper into the garbage. Rinse,

scrub clean, and put away any glassware and the deflagrating spoons. Make sure the bench top is

wiped clean.

DATA & CALCULATIONS 1. Calculate the theoretical yield of oxygen (in g) from the amount of potassium chlorate used.

2. Calculate the actual amount (in g) of oxygen produced using your data.

3. Calculate the % yield.

ANALYZE & APPLY QUESTIONS 1. Write the 6 chemical equations for the reactions that occurred in this experiment.

2. In this experiment, what information did you obtain by testing with the litmus paper?

3. Identify the following reactions as exothermic or endothermic:

a) Pyrolysis of KClO3

b) Reaction of carbon and oxygen

c) Reaction of magnesium and oxygen


to refer back to the purpose. There should be two brief paragraphs—one for Part 1 which involved

quantitative data, and one for Part 2 which was qualitative. In Part 2, use your observations to

describe the physical and chemical properties of oxygen.

41

15. MODEL CONSTRUCTION: POLARITY AND SHAPE

INTRODUCTION The most common type of bond between two atoms is a covalent bond. A covalent bond is

formed when two atoms share a pair of electrons. If both atoms have the same electronegativity,

the BOND is nonpolar covalent. When atoms have different electronegativities, the electrons are

attracted to the atom with the higher electronegativity. The BOND that forms is polar covalent.

MOLECULES made up of covalently bonded atoms may themselves be polar or nonpolar. If

the polar bonds are symmetrical around the central atom, the bonds offset each other and the

molecule is nonpolar. If the polar bonds are not symmetrical, the electrons will be pulled to one end

of the molecule, and the molecule will be polar.

Many physical properties of matter are the result of the shape and polarity of molecule.

Water, for example, has unusual properties that can only be explained by the shape of its molecule

and the distribution of charge on the molecule.

In this lab activity, you will build models of molecules, name their shape and the type of

hybridization, and predict their polarity.

BACKGROUND RESEARCH 1. Prepare a data table for the 20 molecules using the following headings: Formula; Electron dot

structure; Bond type (p/ np); Molecule type (p/ np); Sketch; Shape; Hybridization

2. Draw the electron dot structure for the 20 molecules given.

SAFETY CONCERNS There are no significant safety concerns associated with this lab.

MATERIALS Ball and stick molecular modeling kit

PROCEDURES 1. Complete the table for the following compounds:

1) H2

2) HCl

3) O2

4) N2

5) H2O

6) CO2

7) NH3

8) CH4

9) CH3Cl

10) CCl4

11) C2H6

12) C2H4

13) C2H2

14) HClO

15) HCN

16) H2CO

17) HCOOH

18) H2O2

19) PO43-

20) NO31-

2. a) Construct a model of pentane and draw the structural formula.

b) Construct isomers of pentane and draw structural formulas for each.

3. Construct 2 models of pentene and draw structural formulas.

4. Construct 2 models of dichloropentane and draw structural formulas.

5. Construct a model of bromochlorofluoromethane, and draw the structural formula. Construct an

isomer of this compound. (Hint: Is your left hand identical to your right?)

6. CLEANUP: Completely break down any models. Make sure you can completely close the lid

without forcing it.

42

DATA & CALCULATIONS 1. Complete your data table.

ANALYZE & APPLY QUESTIONS 1. Classify each of the following using one of the following (you can only use each term once): ionic

bonding, polar covalent bonding in a polar molecule; polar covalent bonding in a nonpolar

molecule; nonpolar covalent bonding in a nonpolar molecule

a) Cl2 b) CH2Cl2 c) CCl4 d) KF

2. a) What are isomers?

b) How many isomers were possible for pentane?

3. What type of isomerism (A) Constitutional/ Structural; (B) Cis/ Trans Geometric; (C) Enantiomers/

Diastereomers from chiral carbons) is shown with the various models of the following?

a) Pentane

b) Pentene

c) Dichloropentane

d) Bromochlorofluoromethane

4. Explain why CHCl3 is a polar molecule, but CCl4 is not.

5. Paradichlorobenzene and sodium chloride are both white crystals at room temperature.

Paradichlorobenzene melts at 530C while sodium chloride melts at 12530C. Explain this difference

in melting point.

6. Although methane, water, and ammonia have approximately the same molecular mass, their

melting point and boiling points differ greatly. How do you think molecular shape and polarity

affect boiling and melting point?

7. The polarity of a substance can have a great effect on its soubility. A rough rule of thumb is “like

dissolves like.” Using this general rule, what can you predict about the polarity of ethyl alchol if you

know that alcohol dissolves in water?


43

16. PAPER CHROMATOGRAPHY

INTRODUCTION Chromatography is a technique for separating and identifying components in a mixture. The

name of the method derives from the fact that it was first applied to the separation of colored

substance (Greek: chromos, color0. Although the effectiveness of the separation is most easily

detected if the components of the mixture are colored, the method is applied today to all kinds of

mixtures. All types of chromatography employ two different immiscible phases in contact with each

other. One of the phases is moving, the mobile phase, and the other non-moving phase is the

stationary phase.

In paper chromatography, a small amount of mixture to be separated is placed near the

bottom edge of the paper. This end of the paper strip is then placed in the solvent, which then

travels up the paper by capillary action. Separation occurs because different chemicals in the

mixture travel different distances up the paper due to their attraction to the paper and their

solubility in the solvent. Those substances that are not tightly adsorbed on the paper will move at

the same rate as the solvent. Substances that are bound tightly to the paper will not move at all.

Those substances that are weakly attracted to the paper will move, but more slowly than the

solvent. When the solvent has moved near the end of the paper, the paper is removed from the

solvent and dried. Once developed, the paper called a chromatogram, will contain different

chemicals located at different positions on the paper.

The ratio of the distance each component of the mixture moves with respect to the distance

that the solvent moves can be calculated as follows:

Rf = Dspot where Dspot = distance traveled by the component/ spot

Dfront Dfront = distance traveled by the solvent front

The Rf value is characteristic of a substance under a specified set of conditions and can

sometimes be used in the identification of substance in an unknown mixture when compared to

known pure substances or used to detect the same components in different mixtures.

BACKGROUND RESEARCH 1. In our labs, deionized water is made by pumping tap water through a column packed with beads or

a resin which traps any ions, but lets pure water water flow through. What is the stationary phase

and the mobile phase in this form of column chromatography?

2. What is the stationary and mobile phase in this experiment?

SAFETY CONCERNS There are no significant safety concerns associated with this lab.


44

PROCEDURES 1. Place a small amount of water (to a depth of about 1 cm) in a plastic cup or chromatography vial.

2. Draw a line with a pencil about 2 cm from the edge of the paper. This marks the origin.

3. Put a marker or ink spot on the origin line (don’t use too much ink!), and carefully, keeping the ink

spot above the water, place the paper (spot down) in the cup/ vial containing the solvent (water).

IF YOU PUT THE ORIGIN BELOW THE WATER, YOU MUST REDO YOUR SAMPLE.

4. Do steps 1-3 for all markers or ink you have been given (if you have several people in your group,

you should distribute the work).

5. As the solvent travels up the chromatography paper, make observations.

6. When the solution has nearly reached the top, remove the paper and quickly draw a pencil line

across the paper to mark the distance the solvent has traveled (once dry, you will not otherwise be

able to tell how far it has traveled).

7. Quickly draw a pencil line marking the farthest distance traveled by any component colors as well

(they may continue to bleed upward until the paper is dry, so you want to mark it quickly).

8. Measure the distance traveled from the origin by the solvent for each chromatogram (Dfront), and

measure the distances traveled from the origin by each component/ spot (Dspot) on each

chromatograph.

9. CLEANUP: Pat your chromatograms dry between two pieces of paper towel, then remove and

staple to a dry piece of paper – labeling which chromatogram came from which marker/ ink. Dump

the water in the cup/ vials down the sink, and rinse with tap water. Neatly return all ink/ marker,

chromatography paper, and vials to the location indicated by your instructor. Wipe your tables

clean and dry.

DATA & CALCULATIONS 1. Calculate the Rf value for each component on each chromatogram – indicate the color of the

component.

2. Include a data table which summarizes your measurements. Include the marker/ ink name,

component colors, Ds, Df, and Rfl.

ANALYZE & APPLY QUESTIONS 1. What causes the separation of the ink components to occur?

2. What can you conclude about the ink components that traveled the farthest? That traveled the

least?

3. Which ink(s) appear to be composed of only one compound? Explain, referring to data..

4. Which ink(s) appear to be made up of more than one compound? Explain, referring to data.

5. Do any of the markers appear to contain the same inks? Refer to Rf values.


45

17. DETERMINATION OF ABSOLUTE ZERO

INTRODUCTION Temperature and volume have a directly proportional relationship for gases and can be

represented by a straight line equation in which the independent variable is temperature and the

dependent variable is the volume. When the volume of a fixed amount of gas at a given

temperature is cooled to a lower temperature, the volume decreases. If the volumes and

temperatures at both conditions are accurately measured, these can be plotted. The straight line

that results can be extrapolated to find the point where the line would cross the x-axis (where V =

0). This point should correspond to theoretical absolute zero, since the volume of a gas at absolute

zero should be zero if it has the properties of an ideal gas.

BACKGROUND RESEARCH 1. What variables are being measured?

2. What gas is used in this lab to find absolute zero?

3. What characteristics does an ideal gas have (don’t just mention what you read in the intro)?

4. Which variable will you plot on the y-axis? On the x-axis?

5. What piece of equipment is appropriate for measuring the volume of water in this experiment?

6. When doing Charle’s Law calculations, in what unit must temperature be in?

SAFETY CONCERNS Wear goggles and aprons, and use insulating gloves and tongs to avoid getting burned when handling

hot glassware.


PROCEDURES 1. Set up a boiling water bath by filling a 1-L beaker with tap water and heating on a hot plate (this

may be set up for you already – check your station). While waiting for the water to boil, prepare an

ice water bath by filling a Styrofoam container with ice and tap water. The temperature should be

less than 100C or as close to 00C as possible.

2. Tightly stopper a DRY 250-mL flask with a one-hole stopper and submerge it in the hot water bath

so the air in the flask is completely submerged in the water. THE WATER SHOULD NOT GET INTO

THE FLASK – IF IT DOES, YOU MUST OBTAIN A NEW FLASK AS THIS WILL RUIN YOUR RESULTS!

3. As soon as the water appears to boil, measure the temperature of the water (T1). Do not allow the

thermometer to touch the bottom of the beaker.

4. Leave the flask in the boiling water for at least two to three minutes.

5. Then, remove the flask and quickly invert it into the ice bath – MAKING SURE THAT THE STOPPER

DOES NOT FALL OUT. Push the flask down so it is completely submerged under the ice water—

46

some water will push into the flask (that’s expected and normal!). Allow it to remain there until

the flask has reached the temperature of the water (about 5 minutes). Measure the temperature

of the ice bath when the flask is cold (T2).

6. Meanwhile, TURN OFF/ UNPLUG THE HOT PLATE.

7. When ready to remove, PLACE YOUR FINGERS OVER THE STOPPER HOLE(s) so no more water can

enter or leave the flask. Remove the flask.

8. Measure the volume of water drawn into the flask (Vd).

9. To obtain the volume of the gas at the higher temperature, when it completely occupied the flask,

completely fill the flask with tap water and then insert the rubber stopper. Measure the volume of

water occupying the flask (V1).

10. CLEANUP: Put away an equipment originating from your locker. Empty flask, and place on paper

towel with stopper to air dry on the side of the room. ASK – if you are the last class doing the

experiment for the day, empty the large beakers and Styrofoam containers in your sink, and return

to the side of the room. If another class is coming, you may be asked to leave the hot water and ice

water baths at your station.

DATA & CALCULATIONS 1. Find V2 by calculating the difference between V1 and Vd.

2. Calculate the theoretical V2 using your data and the Charle’s Law equation.

3. Calculate the % error in your V2 using the theoretical V2 as the accepted value.

4. Prepare a graph of Volume vs. Temperature. The temperature values should range from -3000C to

1000C.

5. Find absolute zero (x-intercept, where y = 0) using the equation derived from your graph.

6. Calculate the % error for the experimental value of absolute zero.

7. Include a data table summarizing your measurements, calculations, and percent error.

ANALYZE & APPLY QUESTIONS 1. Explain how this experiment illustrated Charle’s Law.

2. Why would water push into the flask when the hot flask is suddenly submerged into the ice water

bath? Explain why the water was drawn into the flask at the lower temperature.

3. The experiment assumes that at absolute zero, the volume of the gas is zero. Is that correct?

Explain.

4. Would you expect different results if another gas were used in the flask? Explain.


47

18. MOLAR VOLUME OF A GAS

INTRODUCTION The molar volume of any gas is defined as that volume occupied by ONE mole of gas at

standard temperature and pressure (STP). In this laboratory activity, you will be experimentally

determining the molar volume of a gas. Hydrogen gas will be used, since it is easily generated and

collected.

Hydrogen gas can be generated when an active metal reacts with an acid solution. For

example, a specific number of moles of hydrogen can be obtained from reacting a set amount of

magnesium with an excess of hydrochloric acid. The hydrogen gas can be collected by water

displacement in a eudiometer, a gas collection tube calibrated to measure volume to 0.01 mL.

To determine the molar volume, the data must be corrected for laboratory conditions:

• The collected hydrogen gas will be saturated with water vapor; this vapor pressure from the

water must be subtracted from the total pressure to obtain the pressure of dry hydrogen alone.

• The gas will be collected at (non-STP) laboratory conditions. To compare against the accepted

molar volume at STP, an adjustment will need to be made.

• The moles of hydrogen collected will not be equal to one mole, so the molar volume will be

determined by the ratio of the volume of gas collected (at STP conditions) to the number of

moles of hydrogen produced.

BACKGROUND RESEARCH 1. Write the equation for the reaction used to generate the gas in the eudiometer.

2. What two gases will be in the eudiometer?

3. Define STP (be thorough).

4. What information is obtained from a barometer? Why is that needed?

5. What information must be obtained from a reference chart? How will that information be used?

6. Which reactant is limiting?

7. The eudiometer can only measure 50 mL of gas (at most). Assuming a room temperature of 250C

and an atmospheric pressure of 1 atm, what is the maximum mass of magnesium that you would

want to react to produce 30 mL of hydrogen? (You should check that your magnesium is less than

this mass).

SAFETY CONCERNS Wear goggles and aprons. Hydrochloric acid is corrosive


PROCEDURES 1. Mass the strip of magnesium ribbon, and carefully (without breaking it) bend it around the copper

wire on the stopper to secure it in place.

48

2. Fill a 400 mL or 600 mL beaker about 2/3 full with room temperature water.

3. Using the marking on the eudiometer to measure the acid, pour about 8 mL of 6M hydrochloric

acid into the eudiometer.

4. Carefully fill the eudiometer to the top with room temperature water. Since the acid is denser than

pure water, it will remain on the bottom of the tube until the tube is flipped.

5. Place the rubber stopper with the Mg in the tube. There should be no air bubbles in the tube. If

there are, remove the stopper and add more water.

6. Place your finger over the stopper hole and quickly invert the tube into the large beaker of water,

removing your finger once the stopper is under water. Clamp the tube to the ring stand. There is

no danger from the HCl, since it is at the other end of the tube. You will be able to observe the HCl

migrating down the tube and mixing with the water.

7. Observe the reaction. All the magnesium will react in time. While you are waiting for the reaction

to finish, record the atmospheric pressure (in mm Hg) in the room from the barometer (or write it

down, if it is written on the board). Look up the chart for the vapor pressure of water at different

temperatures.

8. When the reaction stops, allow about 5 minutes for the solution cool to room temperature. Then,

measure and record the temperature of the water in the beaker. Assume the collected gas is at the

same temperature. Select the water vapor pressure that is appropriate for the temperature.

9. Gently tap the tube on the bottom of the beaker to release any bubbles of hydrogen on the

bottom. Read the volume of the gas to the nearest 0.05 mL. Be sure to read the bottom of the

meniscus.

10. Empty the collection tube in the waste container, and repeat the experiment if there is time. It is

not necessary to dry the collection tube or change the water in the beaker.

11. CLEANUP: When you are finished, empty contents of the eudiometer and the beakers (which

contain acid) into the waste container. Neatly arrange eudiometer, stopper with copper wire

(make sure you don’t throw this out) at your table. Return beaker, ring stand, and clamp to your

locker. Wipe and dry your bench, and wash your hands.

DATA & CALCULATIONS 1. Complete each of the following calculations for each trial done:

a. Calculate the mole of magnesium used.

b. Calculate the theoretical moles of hydrogen produced using your chemical equation.

c. Determine the partial pressure of the dry hydrogen gas by adjusting for the vapor pressure of

water.

d. Correct the volume to STP conditions. (You can convert to Liters at any point)

e. Calculate the molar volume of hydrogen using the ratio of corrected volume to moles of

hydrogen calculated above.

2. Average the molar volume obtained from both trials (if you did 2 trials).

3. Calculate the % error using the average against the accepted gas molar volume at STP.


49

ANALYZE & APPLY QUESTIONS 1. Consider how the following conditions would affect the volume of hydrogen collected and the

calculated molar volume:

a. Twice as much magnesium was used.

b. The magnesium was not pure (assume other materials are inert).

c. Some hydrogen dissolved in the collecting water.

d. Some air bubbles remained in the eudiometer after water was added.

2. What is the other product? What happens to it?

3. A volume of 35.0 mL of nitrogen gas was collected over water at 230C and 753 mm Hg. The vapor

pressure of water is 21.1 mm Hg.

a. Calculate the partial pressure of the dry nitrogen gas.

b. Calculate the volume of this gas sample at STP.

c. Calculate the experimental volume of one mole of this gas at STP. (Hint: you need to first figure

out how many moles of nitrogen you produced).

CONCLUSION Complete a standard conclusion and error analysis – see “Writing the Lab Report” for help. Your

conclusion should state what you found, and what principles (gas laws) were applied. Make sure to

discuss errors in this experiment, and to describe its impact on your calculation of molar volume. Be

specific depending on whether your results were higher or lower than the accepted value.

50

19. CONSTRUCTION OF AN AIRBAG

INTRODUCTION In modern automobiles, a crash prevention safety feature is the airbag. Contrary to what

most people believe, which is that the gas comes from compressed air tanks, these airbags are filled

with a gas which results from a rapid chemical reaction. You will design your own experiment,

choose which equipment and reagents you will need, and determine the procedure you will use to

construct an airbag. Your airbag for this activity will be a zip-type plastic bag.

The reaction which you will use to generate the gas, CO2, is the reaction between an acid,

HCl, and common household baking soda, NaHCO3. The reaction will generate CO2 gas, sodium

chloride (NaCl), and water (H2O):

NaHCO3 (s) + HCl (aq) � H2O (l) + NaCl (aq) + CO2 (g)

The ideal result is to generate enough gas to fill the bag to plumpness, not to over-inflate or

under inflate the bag. The “airbag” may contain unreacted chemicals and/ or other products of the

reaction.

A note about grading: Your prelab will be reduced to 20 points, and your lab performance

grade will be doubled to 20 points (10 points dependent on the successful inflation of your air bag,

10 points for standard lab performance). Depending on your instructor, instead of writing a formal

lab report, the 60 points normally designated for the writeup may be assessed by evaluation of your

lab manual (materials, procedures, data and calculations) and your typed analyze/ assess questions

and conclusions, so BE THOROUGH AND NEAT!

PURPOSE/ EXPERIMENTAL OBJECTIVES • Design an experiment to solve a problem.

• Use stoichiometry to relate quantities in a chemical reaction.

• Use your knowledge of gas behavior to predict the outcome of a gas generating reaction.

Performance

score for the

successful

inflation of

an air bag

Bag failed to

inflate

Bag inflated

very little or

inflated too

much and

opened

Bag inflated

with a lot of

wrinkles, but

did not open

Bag inflated

without

opening, but

had some

wrinkles

Bag inflated

without

wrinkling or

opening

0 7 8 9 10

BACKGROUND RESEARCH 1. What gas are you generating? How does it interact with the other products made?

2. Consider the other questions in the procedure section before coming to class.

SAFETY CONCERNS Wear goggles and aprons. Hydrochloric acid is corrosive. Inform the teacher of any spills or accidents.

51

MATERIALS Students will be able to use any and all of the following equipment in order to design their experiment.

Any additional supplies must be approved by the teacher. NO additional chemical will be provided.

2 plastic storage bags

0.50 M HCl solution

Lab balances

Thermometers

Funnels

Barometer

Spoons/ scoopulas

Weight boats

Graduated cylinders (10,

25, 50, 100, and 250

mL)

250 mL beakers

Distilled water bottles.

PROCEDURES 1. You will first need to decide how to set up the experiment. In designing your experiment, consider

the following:

• What measurements do you need to make? How will you make them, and what equipment do

you need?

• What calculations will you need to make?

• What quantities of chemicals will you use?

• In what order do you plant to conduct the experiment?

• What safety equipment or personal protective equipment will you use? AND…. What could go

wrong? Plan for and prepare for each worst-case scenario!!

2. Once you have decided what you are going to do as a group, write down a step-by-step procedure

of what you intend to do in your lab notebook (each person must do this) in the left-hand

“Procedure” portion of your lab notebook.

3. Make your material list more specific by detailing the specific equipment, chemicals (with specified

quantities), and other supplies you will be using.

4. In the right-hand side of your lab notebook, show all calculations (with units) that were needed to

determine the amounts of chemicals you will use in your procedure. Leave room for

measurements/ data collected during the performance of your lab.

5. HAVE YOUR TEACHER APPROVE YOUR MATERIALS, PROCEDURES, AND CALCULATIONS BEFORE

CONTINUING. You must all have the initials in your lab notebook.

6. Gather your supplies, chemicals, and equipment.

7. Conduct your experiment. Have your teacher inspect the final results.

8. CLEANUP: When you are finished, clean up your lab area, throw out all trash, return your supplies/

equipment, and wash your hands prior to leaving the lab.

DATA & CALCULATIONS 1. If writing a formal lab report, include a data table summarizing your measurements, calculations. If

not, make sure that everything is neat and organized in your lab manual!

ANALYZE & APPLY QUESTION Airbags must not only produce a consistent and predictable amount of gas, they must also work very

quickly. List at least 4 factors that could increase the speed of your reaction.

52


20. A STUDY OF PHASE CHANGE USING LAURIC ACID

INTRODUCTION Matter can exist in three different physical states—gas, liquid, or solid. In a pure substance,

changes of physical state take place at discrete temperatures which are constant and characteristic

for that substance. The temperature at which a substance changes state from solid to liquid at

standard pressure is called the normal melting point of that substance. The change back to the solid

state occurs at the same temperature and is called the normal freezing point.

In Part 1, you will examine what happens when a pure substance undergoes a change in

physical state, specifically the freezing behavior of an organic compound called lauric acid (C12H24O2),

also known as dodecanoic acid and determine the freezing point.

In Part 2, you will explore the energy change when lauric acid freezes, by determining the

heat released to water in a calorimeter. You will then calculate the heat of fusion (∆Hfus) of the

substance and compare it to the literature value.

BACKGROUND RESEARCH 1. Define the following terms:

a. Calorimeter

b. Heat of fusion

c. Specific heat capacity

2. What is the specific heat capacity of water?

3. What is the literature value for the ∆Hfus for lauric acid? (in J/ mol or J/ g - specify the unit)

4. Is freezing an endothermic or exothermic process? What is happening on the atomic level that

explains this?

SAFETY CONCERNS Wear goggles and aprons. Exercise caution handling the thermometer and the hot plate. To minimize

inhalation of the lauric acid vapors, keep the temperature of the hot water bath below 600C, and do

not allow the lauric acid to go above 500C (if you can help it).


PROCEDURES

PART 1: DETERMINATION OF FREEZING/ MELTING POINT 1. Obtain a test tube containing lauric acid, and remove the stopper.

2. Heat the lauric acid in a beaker of water that has been heated on a hot plate. KEEP THE

TEMPERATURE OF THE WATER BATH BELOW 600C, and DO NOT ALLOW THE WATER TO

CONTAMINATE THE LAURIC ACID. There must be enough water so that the part of the test tube

containing lauric acid is under the water.

53

3. When the lauric acid has just completely melted, place a DRY thermometer in the test tube, making

sure the thermometer is in the center of the lauric acid, and immediately remove the test tube and

thermometer from the hot water. Begin timing and recording temperatures every 30 seconds right

away until the temperature of the lauric acid falls below 400C. One partner should stir and read the

temperature while another times and records. You will probably need to record at least 10

minutes of data, so set up your data table to record the time and temperature.

4. CLEANUP IF STOPPING (If you are doing both parts of the lab in 1 day, then skip this step): If you

are continuing the lab another day, turn off your thermometer (if digital), and place your test tube

with the lauric acid and thermometer in the test tube rack for Part 2. You may be instructed to

leave your hot water baths, but turn off/ unplug your hotplates if another class is coming in, or you

may be asked to empty and put way your beakers and turn off/ unplug your hotplates if you are the

last class doing the lab for the day.

PART 2: DETERMINATION OF HEAT OF FUSION USING CALORIMETRY 1. Obtain a Styrofoam coffee-cup calorimeter with smaller Styrofoam cup lit (with hole for

thermometer). Record the mass of the calorimeter and lid.

2. Place enough water in Styrofoam coffee- cup calorimeter so that it will cover the lauric acid when

the test tube is placed in it, but not so much that the water will splash into the test tube and

contaminate the lauric acid (at least 100 mL of water).

3. Record the mass of the calorimeter, lid, and water.

4. Record the initial temperature of the water (T1).

5. As you did in Part 1 (Step 2), melt the lauric acid, and remove it from the water bath when it has

just melted (DO OT ALLOW THE LAURIC ACID TO GO ABOVE 500C). Remove the thermometer and

clean it with a paper towel. TURN IT OFF (if digital).

6. Let the lauric acid cool in the air while carefully observing for the appearance of crystals. As soon

as the crystals (or any cloudiness) starts to appear, IMMEDIATELY place the test tube in the

calorimeter with lid, and gently swirl. Keep the thermometer reading the temperature of the

water, and record the highest temperature the water reaches (Tf).

7. Remove the test tube. Dry the outside, and record its mass.

8. The number written on the side of the test tube is the mass of the empty test tube. Record this

5. CLEANUP: Replace the stopper in the test tube, and once completely solidified, return the test

tube. You may be instructed to leave your hot water baths, but turn off/ unplug your hotplates if

another class is coming in, or you may be asked to empty and put way your beakers and turn off/

unplug your hotplates if you are the last class doing the lab for the day.

DATA & CALCULATIONS 1. Use data from Part 1 to plot a cooling curve.

2. Calculate the heat of fusion in J/ g for the lauric acid assuming that

Heat lost by lauric acid = - (Heat gained by water)

3. Convert the molar heat of fusion to J/ mol.

4. Calculate the % error for the ∆Hfus.

54


ANALYZE & APPLY QUESTIONS 1. How did you identify your melting/ freezing point using your cooling curve? Explain.

2. During the flat portion of your cooling curve, is your potential energy staying constant or changing?

Is your kinetic energy staying constant or changing?

3. During the diagonal portion of your cooling curve, is your potential energy staying constant or

change? Is your kinetic energy staying constant or changing?

4. In Part 2, what is the effect of the following errors on your Tf and in your calculated value for the

∆Hfus:

a. The lid for the calorimeter was faulty, and some of the heat evolved during freezing was

released to the air rather than to the water.

b. The calorimeter itself absorbed some of the heat evolved during freezing.

c. An overeager lab student place the lauric acid in the calorimeter before it had started to

crystallize/ freeze.

CONCLUSION Complete a standard conclusion and error analysis – see “Writing the Lab Report” for help. The

conclusion should contain 2 paragraphs, one for Part 1 and one for Part 2. Each should include a

statement of your results, a BRIEF discussion of theory and how it was used in this lab. Discuss sources

of error in your experiment for Part 2 ONLY.

55

21. APPLICATION OF SOLUBILITY RULES

INTRODUCTION Solubility can be thought of as the tendency of a substance (the solute) to dissolve in another

substance (the solvent). For qualitative purposes, such terms as “soluble,” “insoluble,” and “slightly

soluble” can be used to describe these tendencies.

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

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

The total number of positive charges is equal to the total number of negative charges in the solution.

If aqueous solutions of two different ionic compounds are mixed, one of the two things will

occur. If all of the ions remain free, no precipitate will form, and the solution will remain clear. If,

however, two oppositely charged ions are attracted to each other strongly enough, they will bond

together to form an ionic compound that is insoluble in water, forming a precipitate.

Using the solubility rules, you will predict whether aqueous solutions will form precipitates

when mixed together, and then carry out the experiment to confirm your predictions. Equations will

be written for those reactions which form precipitates.

BACKGROUND RESEARCH You will receive a photocopy of the data table you will be using in this lab. On the data table in INK, use

the solubility rules to predict whether the solutions will produce a precipitate when mixed. Assume all

are double replacement reactions. Enter yes if a precipitate should form, or no if both products are

soluble.

SAFETY CONCERNS Wear goggles and aprons. Wash your hands when finished.


PROCEDURES 1. Obtain a set of pipets (keep the pipet bulbs down, tip side up to minimize spills and contamination).

2. Add a few (2-3) drops of two reactants (the one on the column with the one on the row). DO NOT

LET THE TIP OF THE PIPET TOUCH, OR CONTAMINATION AND FALSE DATA WILL RESULT.

AgNO3 Na3PO4 KI Ba(NO3)2 Pb(C2H3O2)2 Na2SO4 FeCl3 NaCl

K2CO3 1 9 16 22 27 31 34 36

NaCl 2 10 17 23 28 32 35

FeCl3 3 11 18 24 29 33

Na2SO4 4 12 19 25 30

Pb(C2H3O2)2 5 13 20 26

Ba(NO3)2 6 14 21

KI 7 15

Na3PO4 8

56

3. Use a different color pen to record yes if there is a reaction and no if there is no visible precipitate

(if it is weak, indicate this). Take a picture if you’d like, and make sure you view the mixture over a

dark surface to see if there is any cloudiness.

4. CLEANUP: Replace all solutions, rinse mixtures down sink and wipe your reaction surface off with

soap and running water. Wash hands.

DATA & CALCULATIONS 1. Include a cleaner version of your data table from the reaction of your substances.

2. Include a data table with:

a. Double-replacement equations for any reactions that occurred. Label the precipitate.

b. Net ionic equations for any reactions that occurred.

ANALYZE & APPLY QUESTIONS 1. Use your data to answer the following:

a. Which anions were always spectator ions?

b. List the anions with the total number and formulas of the precipitates each formed.

c. Which cations were always spectator ions?

d. List the cations with the total number and formulas of the precipitates each formed.

2. Compare your results to the predictions based on the solubility rules. Were there cases where a

precipitate formed when it should not have, or when a precipitate did not form when predicted?

In each case, what might have caused the discrepancy?

3. For each of the following pairs of mixture of salts, determine whether the cations could be

separated easily from each other via a precipitation reaction. Explain why or why not. For those

pairs that can be separated, describe what reagent you could use to separate them.

a. Pb(NO3)2 and AgNO3

b. MgCl2 and BaCl2

c. Na2CO3 and K2CO3

d. Li2SO4 and CuSO4


57

22. THE 6-SOLUTION PUZZLE

INTRODUCTION The art of logical thinking is essential to mastering chemistry. Although it is important to be

familiar with many chemical facts so that you can utilize the language of chemistry effectively, it is

even more important to be able to assimilate thee facts and organize them in logical ways so they

can be used to solve problems. The art of logical thinking is not difficult to acquire, but it does take

practice. Try out your own ability on the following logic problem.

In this activity, you will attempt to identify unknown solutions by noting their characteristic

reactions—this is what many forensic tests for various chemical do. You will be given 6 solutions to

react with each other using micro-scale chemistry. The test reactions will be double replacement

reactions that form a precipitate or gas (i.e. fizzing or bubbling). Some solutions have a characteristic

color.

BACKGROUND RESEARCH First, you will warm up by solving a logic puzzle.

During a chemistry class, five pairs of students are working on five different experiments at a row of

lab stations numbered 1-5 consecutively. From the information given below, can you tell what each

student’s experiment is, at which lab station he or she is working, and who his or her lab partner is?

• Hal Ogens (♂) experiment and all the other experiments are performed by lab partners of the

opposite sex. For example, Ann Ion’s (♀) lab partner is a boy.

• Milli Liter (♀) and Ben Zene (♂) work together.

• Charles Law (♂) does NOT work at Station #2.

• Molly Cool (♀) and her partner work between Station #3 and the station occupied by Barry

Um (♂).

• Phyllis Pipet (♀), building a SPECTROSCOPE, is NOT at Station #4.

• Milli Liter (♀) does NOT work next to Molly Cool (♀).

• Earl N. Meyer (♂) works at Station #4.

• IONIC REACTIONS are being carried out at Station #5.

• Hal Ogens (♂) is building a SMALL-SCALE BALANCE.

• Only one of the Um twins, Francie Um (♀) and Barry Um (♂), works at Station #1 doing an

ACID-BASE TITRATION.

• One group is performing a GAS LAW experiment.

SAFETY CONCERNS Wear goggles and aprons. Wash your hands when finished.

MATERIALS Droppers containing 0.1 M solutions of the following: AgNO3, BaCl2, Cu(NO3)2, CuSO4, HCl, Na2CO3

Droppers containing 6 unlabeled solutions.

58

PROCEDURES 1. Obtain a set of droppers, and carry out micro-scale reactions with the known compounds,

recording your observations:

AgNO3 BaCl2 Cu(NO3)2 CuSO4 HCl

Na2CO3 1 6 10 13 15

HCl 2 7 11 14

CuSO4 3 8 12

Cu(NO3)2 4 9

BaCl2 5

2. Repeat the experiments with the unknown compounds, recording your observations.

3. CLEANUP: Replace all solutions, rinse mixtures down sink and wipe your reaction surface off with

soap and running water. Wash hands.

DATA & CALCULATIONS 1. Include a data table summarizing your observations for your known and unknown solutions.

2. Include a data table with:

a. Equations for the reactions that occurred, labeling any precipitates or gases.

b. Identification of your unknown solutions.

ANALYZE & APPLY QUESTIONS 1. What color are copper solutions? 2. What are the three driving forces for a double replacement reaction?

CONCLUSION Complete a standard conclusion and error analysis – see “Writing the Lab Report” for help. Explain

how you identified each unknown solutions – what data were the most useful?

59

23. RATE OF A CHEMICAL REACTION

INTRODUCTION The rate of a chemical reaction is the rate at which reactants are converted to products.

Some reactions are fast and some reactions are slow. The rate of a specific reaction can be found

only by experiment. Reaction rates vary with the temperature and the concentrations of the

reactants.

In this laboratory activity, you will study the effects of concentration of reactants and of

temperature on the rate of reaction between iodate ion (IO31-) and the sulfite ion (SO3

2-). The

reaction is known as the “Iodine Clock reaction,” because a measurable amount of time elapses

before the reaction reaches completion. The reactions that occur are:

(1) IO31- (aq) + SO3

2- (aq) � I1- (aq) + SO42- (aq)

(2) I1- (aq) + IO31- (aq) � I2 (aq) + H2O (l)

(3) I2 (aq) + starch (aq) � dark blue starch-iodine complex

The end of the reaction is signaled by the formation of the colored starch-iodine complex.

In Part, 1, you will run several experiments in which the concentration of the iodate ion is

varied and accurately time the reactions. In Part 2 and 3, your instructor will perform a reaction at

different temperatures and with a catalyst, respectively, and you will then measure the time of

reaction.

Before you begin your experimentation, however, you will need to practice to make sure

your mixing technique is consistent.

BACKGROUND RESEARCH 1. Bring a watch that can time seconds.

2. What are the variables and controls in: a) Part 1, and b) Part 2?

3. What is the total volume when both reactants are mixed in each trial?

4. What is the purpose of the starch in the reaction?

SAFETY CONCERNS Wear goggles and aprons. The solutions contain sulfuric acid, which is corrosive.


PROCEDURES

PRACTICE 1. Obtain solutions A (iodate ion solution) and B (sulfite-starch solution). Clean two 10-mL graduated

cylinders, and label them “A” and “B” respectively so you can measure volumes accurately and

precisely without contamination.

60

2. Measure 3 mL of solution A in one graduated cylinder, and 3 mL of solution B into the other

graduated cylinder. Pour into the individual test tubes, and as you pour the contents of one into

the other, begin timing. Mix by pouring the contents back and forth 3-4 times, and note the time

as soon as a color appears. Record it!

3. CLEANUP: Dump contents into waste beaker, and rinse test tubes and shake dry (it’s okay, if the

test tubes are not completely dry). It is NOT necessary to rinse graduate cylinders.

4. Repeat this procedure until you get two close readings (within 2 seconds).

PART 1: EFFECT OF CONCENTRATION ON TIME 1. Prepare different concentrations of solution A as shown in the data table below. During your first

trial, record the temperature of solution A and B (with separately labeled thermometers that have

been rinsed and cleaned before use).

Add the water to solution A in the graduated cylinder so that the volume with each dilution is 7 mL.

Proceed with reacting the solutions as you practiced before for each dilution, performing 2 trials

for each dilution.

Trial Vol A (mL) Vol H2O (mL) Vol B (mL) Time (s), Trial 1 Time (s), Trial 2

1 6 1 3

2 5 2 3

3 4 3 3

4 3 4 3

5 2 5 3

6 1 6 3

2. Trouble-shooting: If the times for your 2 trials are more than 5 seconds apart, perform a third trial,

and choose the two closest measurements.

3. CLEANUP: Cleanup as you did when you were practicing. Notify your instructor before the waste

beaker becomes too full.

PART 2 & 3 : EFFECT OF TEMPERATURE & A CATALYST ON TIME (Demo) 1. The concentrations used in trial 3 will be run at two additional temperatures (one hotter, one

cooler). Record the time. Two trials will be done.

2. The concentrations used in trial 3 will be run at room temperature, but Cu+2 (added by the addition

of some Cu(NO3)2 solution) will catalyze the reaction. Two trials will be done. Record the time.

3. CLEANUP: The solutions will be dumped in a waste beaker for treatment, and glassware cleaned.

DATA & CALCULATIONS 1. Calculate the [IO3

1-] for each dilution if your stock solution of A, (IO31-) was made with 4.3 g of KIO3

for 1 liter of solution. 2. Calculate the average time in seconds for each trial. 3. Calculate the rate (1/s) for each trial by taking the inverse of each average time. 4. Include a data table summarizing concentration, average time, and rate for each trial. 5. Prepare the following graphs:

a. Time vs. [IO31-]

61

b. Rate vs [IO31-]

6. Calculate the slope from the rate vs. [IO31-].

ANALYZE & APPLY QUESTIONS 1. What does the slope represent in your second graph?

2. What is the order of the reaction with respect to IO31-? Write the rate law expression AS

DETERMINED BY DATA IN THIS LAB.

3. Can the overall order of the reaction be determine from this experiment alone? Explain.

4. How should the temperature affect the rate? Was this effect observed?


to explain your graphical analysis.

62

24. A STUDY OF CHEMICAL EQUILIBRIUM

INTRODUCTION Le Chatelier’s principle says that if a reaction system at equilibrium is stressed, the reaction

will shift in a direction that will relieve the stress (either forward or backward). For example, if we

add a reactant, the reaction will shift toward the right to use some of the reactants to make products

and reestablish a balance of reactants and products that is in equilibrium. Similarly, if we add

products, the reaction will shift toward the left to use some of the products to make reactants and

reestablish a balance of reactants and products that is in equilibrium. Keep in mind that BOTH

REACTANTS AND PRODUCTS ARE PRESENT IN THE REACTION TEST TUBES.

Color is an easily observed macroscopic property that can be used to indicate shifts in

equilibrium concentrations. In this lab, you will investigate chemical systems that are already at

equilibrium. You will disturb these systems by adding or subtracting reactants or products and

observe color changes that indicate how the equilibrium systems respond. You will then explain

those changes in terms of Le Chatelier’s principle.

BACKGROUND RESEARCH 1. When a reaction shifts forward/ right, what happens to the amount of reactants as a result? What

happens to the amount of products as a result?

2. When a reaction shifts backward/ left, what happens to the amount of reactants as a result? What

happens to the amount of products as a result?

3. Write the equilibrium expressions (mass action expressions) for each reaction that you will be

working with (remember… you need to know the physical state of substances – s, l, g, aq, to do this

properly).

4. What is the solubility of Fe(OH)3? (This should help you analyze System 2).

5. Make sure to include the given data tables for each section into your lab book:

Stress Species

added or

removed

Change in concentration (or amount) as a result

of reaction shift: Indicate INC, DEC, or N.C. (no

change)

Observations

System 1

Stress 1

Stress 2

Stress 3

X

____

____

____

Yellow Blue

HBB (aq) � H+1(aq) + BB-1

(aq)

___________________

___________________

___________________

___________________

____

____

____

____

____

____

____

____

____

System 2

Stress 1

Stress 2

Stress 3

X

____

____

____

Pale brown Dark red

Fe+3(aq) + SCN1-

(aq) � FeSCN+2

(aq)

___________________

___________________

___________________

___________________

____

____

____

____

____

____

____

____

____

System 3

Stress

X

____

NaCl (s) � Na+(aq) + Cl1-

(aq)+

____ _____ ___

___________________

___________________

63

System 4

Stress 1

Stress 2

Stress 3

Stress 4

Stress 5

Stress 6

X

____

____

____

____

____

____

Pink Blue

Co(H2O)6+2

(aq) + 4 Cl-(aq)� CoCl42-

(aq)+ 6 H2O (l)

___________________

___________________

___________________

___________________

___________________

___________________

___________________

____

____

____

____

____

____

____

____

____

____

____

____

____

____

____

____

____

____

____

____

____

____

____

____

SAFETY CONCERNS Wear goggles and aprons. HCl and NaOH are corrosive and can cause burns, so avoid skin contact.

KSCN (potassium thiocyanate) is poisonous – use with care. CoCl2 is a carcinogen – minimize exposure,

and do not use more than needed.


PROCEDURES

SYSTEM 1: ACID-BASE EQUILIBRIUM 1. SYSTEM: Add 4 drops of bromothymol blue (HBB) to about 10 mL of distilled water – the system

will already be at equilibrium. Pour half of it into another test tube as a control that will remain

unreacted for comparison. Observe the system. Perform the stresses to the other test tube.

2. STRESS 1: Add 0.1 M HCl a drop at a time until a color change occurs. Observe the system,

compare to the control.

3. STRESS 2: Add 0.1 M NaOH a drop at a time until a color change occurs. Observe the system,

compare to the control. While OH- is not part of the equilibrium system, OH- reacts with H+.

4. STRESS 3: Now add 0.1 M HCl a drop at a time until a color change occurs again. Observe the

system, compare to the control.

5. CLEANUP: Dump the contents into the acid-waste beaker, and clean the test tubes with tap water

and a test tube brush.

SYSTEM 2: IRON (III) THIOCYANATE EQUILIBRIUM 1. SYSTEM: Get 4 separate test tubes, and place about 4 mL iron (III) thiocyanate in each test tubes.

Observe the system, noting the color. Set aside the first test tube as a control used for comparison.

2. STRESS 1: To the second test tube, add a few crystals of Fe(NO3)3. Observe the system, compare to

the control.

3. STRESS 2: To the third test tube, add 10-15 drops of KSCN solution. Observe the system, compare

to the control.

4. STRESS 3: To the fourth test tube, add a few drops of 6 M NaOH. Observe the system, compare to

the control.

5. CLEANUP: Dump the contents into a designated THIOCYANATE waste beaker. Clean your test

tubes with tap water and a test tube brush.

64

SYSTEM 3: SODIUM CHLORIDE SOLUBILITY EQUILIBRIUM 1. SYSTEM: In a clean test tube, obtain about 5 mL of saturated NaCl solution.

2. STRESS 1: Add a few drops of concentrated HCl. Observe the system, and record results.

3. CLEANUP: Dump the contents into the acid-waste beaker, and clean the test tubes with tap water

and a test tube brush.

SYSTEM 4: COBALT (II) COMPLEX IONS 1. SYSTEM: Place a very small crystal of CoCl2 in a small test tube. Add about 1 mL distilled water until

you make a purplish solution (not pink or blue). Your instructor may already have made this

solution for you. Pour half of your solution into another test tube as a control that will remain

unreacted for comparison. Observe the system. Perform the stresses to the other test tube.

2. STRESS 1: Place the test tube in a hot water bath. Observe the system, compare to the control.

3. STRESS 2: Place the test tube in an ice water bath. Observe the system, compare to the control.

Allow test tube to come to room temperature before continuing.

4. STRESS 3: Add acetone until the color changes (Acetone removes some of the water). Observe the

system, compare to the control.

5. STRESS 4: Add water until the color changes. Observe the system, compare to the control.

6. STRESS 5: Add solid CaCl2, shaking to dissolve . Continue adding and dissolving more CaCl2 until

there is a color change. Observe the system, compare to the control.

7. STRESS 6: Add AgNO3 solution. Observe the system, compare to the control.

8. CLEANUP: Dump the contents into a designated COBALT waste beaker. Clean your test tubes with

tap water and a test tube brush.

DATA & CALCULATIONS 1. Complete your data table.

ANALYZE & APPLY QUESTIONS 1. What is the most common chemical method for removing H+ (H3O+) ions in aqueous solution?

Write a net ionic equation which describes this method.

2. In system 4, on which side of the chemical reaction should the “heat” term be placed?

3. Predict the shift (to REACTANTS or PRODUCTS) and ANY OBSERVABLE VISIBLE CHANGE if the

following were added to your equilibrium systems (answer in a data table format):

a. System 1 i) H2SO4 ii) KOH solution

b. System 2: i) Ca(SCN)2 solution ii) K2CO3 solution iii) NaBr solution

iv) Ca(OH)2 solution v) NaNO3 solution

c. System 3: i) NaNO3 solution ii) NaCl solid

d. System 4: i) Pb(NO3)2 solution ii) NaBr solution iii) NaCl solution

CONCLUSION Complete a standard conclusion and error analysis – see “Writing the Lab Report” for help. Explain

LeChatelier’s principle. Use TWO SYSTEMS from the lab as examples to illustrate it. Note which

direction the reaction shifts.

65

25. CHANGE IN ENTHALPY OF A REACTION

INTRODUCTION When a chemical reaction takes place, chemical bonds in the reactants are broken and new

chemical bonds in the products are formed. Energy is always absorbed in the breaking of bonds and

always released when bonds are formed. If the energy required to break old bonds is less than the

energy released in forming new bonds, the difference in energy is given off and the reaction is said

to be exothermic. The enthalpy change for an exothermic reaction is given a negative sign to

indicate that energy flows from the system.

In this experiment, you will measure the amount of energy released by the decomposition of

hydrogen peroxide, H2O2, into water and oxygen and compare the experimentally determined result

with a calculated value.

The common disinfectant, 3% hydrogen peroxide (H2O2) solution, will be used with

manganese dioxide as a catalyst for the decomposition. The amount of energy released in the

reaction can be calculated from the mass of the water in the solution, the change in temperature,

and the specific heat of water.

BACKGROUND RESEARCH 1. Write the balanced equation for the decomposition of hydrogen peroxide into water and oxygen

gas.

2. Use Hess’ Law to calculate the net energy released. Then calculate the energy in kJ/ mol—this will

be the theoretical value:

H2 + ½ O2 � H2O ∆H = -286 kJ/ mol

H2 + O2 � H2O2 ∆H = -191 kJ/ mol

SAFETY CONCERNS Goggles must be worn at all times.


PROCEDURES 1. Obtain a Styrofoam coffee cup calorimeter with lid, and place two small scoops of manganese

dioxide into the calorimeter.

2. Measure about 45 mL of 3% hydrogen peroxide, and record the exact volume.

3. Measure the temperature of the hydrogen peroxide in the graduated cylinder, and record it.

4. Pour the hydrogen peroxide into the calorimeter, immediately covering with the lid and pushing

the thermometer through the lid. Swirl the calorimeter gently and not the temperature change,

recording the highest temperature.

5. Open the calorimeter, and look inside. Record your observation.

6. Empty the calorimeter into the waste beaker, and rinse it out.

7. Repeat the experiment at least once more. The results will be shared with the class.

66

8. CLEANUP: Empty the calorimeter into the waste beaker. Rinse it out, and return it. Turn off and

return thermometer, and put away glassware.

DATA & CALCULATIONS 1. Number the following calculations and make sure to show setup:

1) Calculate the change in enthalpy for the amount of H2O2 in your data using q = mC∆T. (You

may assume that the density of H2O2 is the same as water).

2) Convert from joules to kilojoules.

3) Calculate the mass of H2O2 in the solution of 3% hydrogen peroxide.

4) Calculate the moles of H2O2.

5) Calculate the ∆H in kJ/ mol.

6) Calculate what % of the theoretical ∆H you were able to measure in this experiment. (Refer to

your background research for the theoretical value).

7) The values for the class will be collected and analyzed. Calculate a class average, and use this

average to find the class % error.

2. Include a data table summarizing your data and calculations, including your percent error.

ANALYZE & APPLY QUESTIONS 1. How will the mass of the system change throughout the experiment? Explain.

2. Compare your results to the theoretical value. How do they compare? Refer to your data in your

answer.

3. Suggest two reasons for the difference between your results and the theoretical value.

4. Although the reaction is exothermic, explain why it doesn’t feel warm when hydrogen peroxide is

put on a cut.

5. The decomposition of hydrogen peroxide is slow. If poured on the table, there is no evidence of

decomposition. On the other hand, when 3% hydrogen peroxide is applied to a cut, the

decomposition begins immediately (bubbles appear). Use your knowledge of biology to suggest a

reason why this occurs.


67

26. DETERMINATION OF pH USING INDICATORS

INTRODUCTION

BACKGROUND RESEARCH 1.

SAFETY CONCERNS Goggles and aprons must be worn at all times.


PROCEDURES 1.

DATA & CALCULATIONS 1.

ANALYZE & APPLY QUESTIONS 1.


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