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Academic/Honors Chemistry Laboratory Guide

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IntroductionWelcome to chemistry here at WHS. Ms. Cipolla and Mr. Rickard are looking

forward to helping you get the most out of this class during the course of the year. One of the special things about Chemistry, other than the teachers themselves, is that it is a laboratory class that meets six times each week. Once a week, in addition to your regularly scheduled class time, you will have class two periods in a row. This extra period is intended to provide an extended period of time for activities like experiments that take longer than a single period.

While we tend to think of it as a “lab period,” experiments don’t always happen during the double period day. Experiments can be and are performed as appropriate in the class schedule. Ms. Cipolla and Mr. Rickard recommend always having a pair of closed toed shoes in your locker just in case, even if it isn’t your class’s day for a double period.

This manual is meant to be your personal trainer and guide while working in the laboratory. Sometimes you will have a simple question two which you need an answer or maybe you just need a reminder of how to setup or use a particular piece of equipment. This guide can help you with those issues so that you do not need to stop your experiment until your teacher can answer your question. If you follow the procedures in this manual, you will be more likely to get better results and you will score more points as well.

If you have suggestions for things that should go in the lab manual or if you find typos or other mistakes as we work through the year, please bring them to Mr. Rickard’s or Ms. Cipolla’s attention so that the next version of the manual can be even better.

Ms. Cipolla & Mr. Rickard

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Table of Contents

ACADEMIC/HONORS CHEMISTRY LABORATORY GUIDE..........................................................................1

INTRODUCTION..................................................................................................................................... 3

TABLE OF CONTENTS.............................................................................................................................. I

LAB SAFETY........................................................................................................................................... 3

LABORATORY RULES.......................................................................................................................................3LABORATORY DRESS CODE..............................................................................................................................5

LAB EQUIPMENT.................................................................................................................................... 6

SAFETY TEST REVIEW GUIDE................................................................................................................ 17

KEEPING A LABORATORY NOTEBOOK (HONORS CHEMISTRY ONLY).....................................................18

RULES FOR WRITING IN YOUR LABORATORY NOTEBOOK.......................................................................................18SETTING UP YOUR LABORATORY NOTEBOOK......................................................................................................18ADDING AN EXPERIMENT (PRE-LAB ASSIGNMENT).............................................................................................19DURING THE EXPERIMENT..............................................................................................................................19AFTER THE EXPERIMENT................................................................................................................................20EXAMPLE NOTEBOOK ENTRY FROM MR. RICKARD’S NOTEBOOK.............................................................................20

LABORATORY PROCEDURES................................................................................................................. 23

HOW TO USE A BUNSEN BURNER....................................................................................................................23HOW TO USE A HOT PLATE.............................................................................................................................23HOW TO USE FILTER PAPER FOR GRAVITY FILTRATION..........................................................................................24HOW TO TAKE MEASUREMENTS......................................................................................................................25

LABORATORY EXPERIMENTS................................................................................................................ 28

EXPERIMENT TITLE: GOLDEN PENNY................................................................................................................29EXPERIMENT TITLE: CANDLE OBSERVATIONS.....................................................................................................30EXPERIMENT TITLE: SEPARATING THE SLUDGE..................................................................................................33EXPERIMENT TITLE: MEASUREMENT MAKING....................................................................................................37EXPERIMENT TITLE: DENSITY..........................................................................................................................40EXPERIMENT TITLE: HYDRATES.......................................................................................................................42EXPERIMENT TITLE: PRECIPITATION.................................................................................................................45EXPERIMENT TITLE: SPECIFIC HEAT..................................................................................................................47EXPERIMENT TITLE: CALORIMETRY – HEAT OF FUSION OF ICE..............................................................................49EXPERIMENT TITLE: MOLAR VOLUME OF HYDROGEN GAS...................................................................................52EXPERIMENT TITLE: PH.................................................................................................................................54EXPERIMENT TITLE: CANDY COATING CHEMISTRY..............................................................................................55EXPERIMENT TITLE: STRONG ACID AND STRONG BASE TITRATION.........................................................................58THE GOLDEN FLEECE....................................................................................................................................59WHITE POWDERS LAB..................................................................................................................................60CHROMATOGRAPHY.....................................................................................................................................63

MEASURING PLASTICS ...............................................................................................64PERCENT SUGAR IN BUBBLE GUM NAMES:...........................................................................67

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MATERIALS:......................................................................................................................................... 67

CHALK IT UP TO CHEMISTRY!..........................................................................................................................69THE MOLAR VOLUME OF HYDROGEN GAS........................................................................................................72INTRODUCTION/PURPOSE..............................................................................................................................72PRECAUTIONS.............................................................................................................................................72WEAR GOGGLES. HYDROCHLORIC ACID IS CORROSIVE........................................................................................72PROCEDURE................................................................................................................................................72CALCULATIONS- SHOW THE SET UP FOR EACH OF THE FOLLOWING.........................................................................73HYDRATE LAB.............................................................................................................................................74PENNY ISOTOPES LAB...................................................................................................................................76CLASSIFICATION OF CHEMICAL SUBSTANCES......................................................................................................78CHEMICAL REACTIONS LAB............................................................................................................................80STRUCTURE OF COMPOUNDS PART 1..............................................................................................................82STRUCTURE OF COMPOUNDS PART 2..............................................................................................................85DENSITY OF LIQUIDS.......................................................................................................................................2PHYSICAL & CHEMICAL CHANGES.....................................................................................................................5DENSITY OF CARBON DIOXIDE..........................................................................................................................9FLAME TESTS..............................................................................................................................................11METALS VS. NONMETALS..............................................................................................................................13EXPERIMENT TEMPLATE................................................................................................................................15Experiment Title: INSERT EXPERIMENT TITLE HERE................................................................................15

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Lab SafetyWhile it may not be the most glamorous topic in chemistry, laboratory safety is a serious matter.

One of the things students say they are most looking forward to in chemistry is “blowing stuff up.” May of the experiments we do involve things like dangerous chemicals, fire, and fragile glassware. Knowing how to be safe is the key to being able to do some of the more difficult and exciting experiments.

Laboratory Rules

General Rules1. All students must read the Laboratory Rules and Dress Code and return the signed block to Your instructor

prior to being allowed into the laboratory area.

2. Students are permitted into the laboratory area only when given teacher permission.

3. All students must follow the dress code detailed below to be permitted into the laboratory area.

4. Medical conditions such as epilepsy, pregnancy, and dyslexia can be particularly dangerous in the laboratory. We understand that many medical conditions are personal in nature. Please make sure that the school nurse is aware of any ongoing medical conditions or medical conditions that may develop during the course of the year. If you are uncomfortable with Your instructor knowing the nature of your medical condition, the nurse will provide your teacher with guidelines as to what is or is not acceptable for you in the laboratory. Remember, chemicals don’t know what medical conditions you have and can also cause additional problems.

Behavior1. The classroom is a chemical laboratory. Therefore, no food or drink is permitted unless it is supplied for

an experiment.

2. Students may only perform the experiment assigned as instructed. Unauthorized changes to the experiments (regardless of how interesting they may seem) are not permitted.

3. Wash your hands when the experiment is finished or when leaving the lab for any reason to prevent contamination. (Chemists are known for washing their hands BEFORE they go to the bathroom.)

4. If you are not feeling well, report it to Your instructor as soon as possible, especially if you believe it is a result of chemical exposure.

5. Never smell a chemical directly from the original container. Use the wafting technique when smelling a chemical is absolutely necessary.

6. Students are not required to work with chemicals originating from unlabeled containers except when issued an unknown to identify.

7. Read the labels on chemical containers before you use them.

8. Estimate how much chemical you need and take only what is necessary.

9. NEVER return unused chemicals to the original container (see the above rule) because this can cause contamination and result in the entire bottle being unusable. First see if any groups around you need that chemical and, if not, return the unused chemical to Your instructor.

10. NEVER pour chemicals down the drain or place them in the garbage can without approval from Your instructor.

11. Keep your lab bench as neat as possible at all times.

Equipment1. Students must know where safety items such as MSDSs, Eye Washes, Safety Shower, and Fire Blanket are

located and how to use them.

2. Report any accidents, spills, and broken glassware to Your instructor immediately.

3. Broken glassware should not be picked up by hand. Broken glassware should be cleaned up with a dustpan and broom when possible. If glass pieces cannot be collected using a dustpan and broom, a damp paper towel should be used.

4. Do not use chipped or cracked glassware. Take them to Your instructor to obtain a replacement.

5. Hot objects in lab often stay hot for an extended period of time. Mr. Rickard has a great hot plate story from graduate school that illustrates this point.

6. NEVER look down the opening of ANY container such as beakers, flasks, or test tubes containing chemicals as a splatter or explosion could occur. Glass is clear for a reason.

7. Do not use graduated cylinders for purposes other than measuring liquids. Storing chemicals and reacting of chemicals should be done in beakers, flasks, and test tubes as instructed.

8. Thermometers are not a substitute for stirring rods.

Fire1. If an unplanned fire occurs, DO NOT PANIC and tell Your instructor immediately. Turn off all gas taps as

long as it is safe to do so.

2. Never move an object that is burning to avoid spreading the fire.

3. Do not put water on a fire unless instructed. Chemical fires play by their own rules.

4. If the fire can be covered with the fire blanket or a large beaker to cut off the oxygen supply, this is a good first idea.

5. If the fire is too large to cover, Your instructor will use the fire extinguisher when appropriate.

6. If a part of your clothing or hair is on fire, use the fire blanket or the safety shower as instructed.

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Laboratory Dress Code1. Safety goggles will be provided and must be worn in a proper manner at all times in the laboratory area.

Splashed chemicals are the most common cause of injury other than hot objects causing burns. (Goggles are meant to protect eyes, not foreheads and necks)

2. Laboratory aprons must be worn when chemicals are being used. Aprons are chemical and flame resistant and can help save you and your clothing.

3. All students must wear closed toe shoes. Sandals and open toe shoes are not permitted. (If you think you will forget, leave a pair of shoes in your locker for lab days.)

4. All long hair and bangs must be tied back when Bunsen burners are being used. (Excessive hair spray is also not recommended.)

5. Loose clothing is not recommended when working with Bunsen burners because it becomes a fire hazard. It is also not recommended because it can knock over glassware or absorb liquid chemicals.

6. Exposed skin around the waist is not permitted because corrosive chemicals are used in the laboratory and may come into contact with skin when leaning against benches.

7. Wearing jewelry in the lab is not recommended. Jewelry can react with chemicals and can also trap chemicals against your skin. Dangling jewelry can also present the same hazards as loose clothing. Watches are permitted but it is recommended that you remove them before washing your hands at the end of the experiment and then rinse the watch as well.

Lab Equipment

Apron: An Apron is one of the most common forms of protective equipment in the laboratory. It helps to protect against chemical spills and is also somewhat fire resistant.

Balance, Electronic:

An Electronic Balance is used to measure the mass of chemicals and equipment in the laboratory. Balances have a variety of ranges and levels of precision.

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Beaker: A Beaker is used to hold and heat liquids. Multipurpose and essential in the lab.

Bottle: Bottles can be used for storage, for mixing, and for displaying.

Bunsen Burner:

Bunsen burners are used for heating and exposing items to flame. They have many more uses than a hot plate, but do not replace a hot plate.

Burette:aka Buret

The Burette is used in titrations to measure precisely how much liquid is used when a high level of accuracy is needed.

Clamp, Burette:

A Burette Clamp is used for holding burettes while in use in the laboratory. Most burette clamps can hold two burettes at a time. Take care when using these clamps because they have a tendency to tip ring stands if not used correctly.

Clamp, Extension:

An Extension Clamp is used for holding a variety of objects in lab. Be careful not to uses clamps with plastic coating for objects that are going to be heated to high temperatures.

Clamp Holder:

A Clamp Holder can be used to hold clamps and metal rods in the laboratory.

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Crucible: Crucibles are used to heat small quantities to very high temperatures.

Erlenmeyer Flask:

The Erlenmeyer Flask is used to heat and store liquids. The advantage to the Erlenmeyer Flask is that the bottom is wider than the top so it will heat quicker because of the greater surface area exposed to the heat.

Evaporating Dish:

The Evaporating Dish is used to heat and evaporate liquids.

Funnel: The Funnel is a piece of equipment that is used in the lab but is not confined to the lab. The funnel can be used to target liquids into any container so they will not be lost or spilled.

Gas Collecting Bottle:

The Gas Collecting Bottle is designed to collect gases produced when used with a water bath. They may also be used like a normal bottle.

Goggles: Goggles are some of the most common pieces of protective equipment. They are used to protect eyes from spills, splashes, and projectiles. Goggles we will use will have splash guards across the top as well as a strap for adjusting the fit.

Graduated Cylinder

Graduated Cylinders are used to measure volumes of liquids when a medium level of accuracy is needed. They come in a variety of sizes.

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Hot Plate: Hot Plates are uses to heat objects and chemicals in the laboratory. They have adjustable power controls. Due to the ceramic top, they will remain hot long after the power has been turned off.

Mortar and Pestle:

The Mortar and Pestle are used to crush solids into the powders for experiments, usually to better dissolve the solids.

Paper Towels:

Paper Towels are essential to the lab environment. They will be used in almost every lab

Pipet: The Pipet is used for moving small amounts of liquid from place to place. Some pipets come with graduation lines for a high level of accuracy.

Pipet Bulb: A rubber Bulb used with a pipet to provide the suction needed to draw liquid chemicals into the pipet. They are designed to eliminate the need for the dangerous method of pipeting by mouth. Some bulbs come with valves for added control.

Ring Stand: Ring Stands are used to hold items being heated. Clamps or rings can be used so that items may be placed above the lab table for heating by Bunsen burners or other items.

Ring Support:

Ring Supports are used with other pieces of equipment such as Triangles and Wire Mesh to support other pieces of equipment as needed.

Rubber Tubing:

Rubber Tubing is used for a variety of things in lab including transferring chemicals and gases during an experiment. It should be checked for cracks before use.

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Scoopula: A Scoopula is a metal scoop used to remove chemicals from a reagent bottle.

Stir Rod: The stir rods are used to stir things. They are usually made of glass. Stir Rods are very useful in the lab setting.

Stopper: Stoppers come in many different sizes. The sizes are from 0-14. Stoppers can have holes for thermometers and for other probes that may be used.

Test Tube: The test tube is the vessel of choice for many reactions and for heating small amounts of chemicals over a Bunsen burner or in a water bath.

Test Tube Brush:

The test tube brush is used to easily clean the inside of a test tube.

Test Tube Holder:

The Holder is used to hold test tubes when they are hot and untouchable.

Test Tube Rack:

The test tube rack is used to hold test tubes while reactions happen in them or while they are not needed.

Thermometer:

The thermometer is used to take temperature of solids, liquids, and gases. They are usually in ºC, but can also be in ºF. All thermometers we will use contain alcohol and not mercury for safety purposes.

Tongs: Tongs are used to hold many different things such as flasks, crucibles, and evaporating dishes when they are hot.

Triangle: The triangle is used to hold crucibles when they are being heated. They usually sit on a ring stand.

Tweezers:aka forceps

Tweezers are used to handle small objects in the laboratory such as many crystalline chemicals.

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Volumetric Flask:

The Volumetric Flask is used to measure one specific volume. They are mostly used in mixing solutions where a one liter or one half liter is needed.

Washing Bottle:

A Washing Bottle contains water and is used to help clean glassware such as test tubes and flasks. It is not a squirt gun. When squirt guns are required in lab, they will be supplied by Mr. Rickard.

Watch Glass:

The watch glass is used to hold solids when being weighed or transported. They should never be heated.

Weigh Paper:

Weigh Papers are used to protect the balance pan from chemicals and to make it easy to handle small amounts of powder chemicals.

Well Plate: Well Plates are used to temporarily store chemicals for reactions and are also used as a place to perform reactions on a small scale.

Wire Gauze: Wire Gauze is used to support items such as beakers and crucibles on a ring stand. It may also be used like a trivet on the lab bench to help hot objects to cool down. Wire Gauze works much like Wire Mesh but also aids in heat dispersion during both heating and cooling.

Wire Mesh: Wire Mesh is used to support items such as beakers and crucibles on a ring stand. It may also be used like a trivet on the lab bench to help hot objects to cool safely.

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Safety Test Review Guide

The following topics are things you should know for the safety test. For safety items found in the laboratory, you should also know where they are located. You will not be permitted to enter the laboratory until you have passed this test. It is not meant to be hard. It is meant to keep everyone safe and injury free throughout the year.

Safety Equipment Location and Proper Use including fire extinguishers, fire blanket, first aid kit, eyewash station, safety shower, goggles, aprons, etc.

Safety Vocabulary including but not limited to carcinogen, mutagen, teratogen, flammable, inflammable, inhalation, ingestion, injection, absorption, corrosive, oxidizer, MSDS, acute effect, chronic effect, irritant, caustic, toxic

Laboratory Dress Code Safety Procedures Glassware and Laboratory Equipment both descriptions and proper uses

Keeping a Laboratory Notebook (Honors Chemistry only)

A laboratory notebook is a common tool found in a majority of scientific laboratories in one way, shape, or form. Traditionally, laboratory notebooks have been bound notebooks that can survive the rigors and abuses of being in the laboratory while experiments are being performed. Many modern laboratories are taking advantage of computers and using electronic laboratory notebooks that exist on the computer network and can be accessed by people in a variety of locations.

Regardless of type, the point of a laboratory notebook is to have a record of what was performed in the laboratory, how it was performed, what results were observed, show how calculations were completed, record thoughts and ideas of the scientist, and at times to explain those ideas and the results obtained. The general rule of thumb for a laboratory notebook is that someone else should be able to pick up the notebook and using the information in the notebook, recreate the experiment, and hopefully get a similar result.

In honors chemistry this year we will use a traditional, dead tree notebook. You will use your notebook to help prepare for experiments before you will be allowed to perform the experiment in the laboratory. It will be your record of what you did in the laboratory and it will be the major source of the laboratory grades for the course. Develop good notebook habits now and they will carry you through the rest of your science classes and beyond.

Rules for writing in your laboratory notebook1. Use only PERMANENT BLACK or BLUE INK in your laboratory notebook. Pencil, erasable

ink, crayon, and other colors of ink are not acceptable and will result in a deduction when the notebooks are graded.

2. Write only on the right side pages. Left side pages may not be written on by students. Graphs, charts, and other visual aids may be glued to the left hand pages as instructed. Leaving the left hand page blank also leaves your instructor a place to write comments and scores when they grade the experiment.

3. If you make a mistake, draw a SINGLE LINE through the mistake so that the mistake may be seen and understood but will be recognized as an error and will be ignored. You may then continue writing the correct information in the laboratory notebook. NEVER scribble out the mistake or try to erase the mistake. Scribbles and erasures will result in a deduction when the notebooks are graded.Example: When you make a misteak in UR mistake in your notebook, simply strike it with a single line and then continue.

4. No white out.5. Do not write in the page margins or in the half space below the bottom line.6. Start a new page for each experiment.7. Please write legibly. If a laboratory notebook cannot be read, the point of having a notebook in

the first place has been utterly defeated.8. When explaining, summarizing, or answering use complete sentences. Complete sentences are not

necessary when recording data or completing calculations.

Setting up your laboratory notebook1. Obtain a bound composition notebook in which the pages are bound and not easy to remove.2. Write your name, course name, and period on the front cover.3. Write your name and homeroom on the inside cover.

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4. Write and underline Table of Contents on the top line of the first page. (HINT: Rulers are very helpful for drawing straight lines.)

5. Skip one line and write and underline Experiment Title and on the same line you also will add small columns for Page and Grade with those titles underlined as well. Each time you perform an experiment, you will record its title in this list as well as the page number in the notebook where the experiment begins.

6. In the lower right corner of the right side pages, number each page in the notebook. Remember, you will not be using the left side pages for writing. Take your time with this step to ensure none of the pages have stuck together.

Adding an experiment (Pre-Lab Assignment)Prior to an experiment, you will usually be given an assignment to prepare your laboratory

notebook for the experiment you are about to accomplish. This must be completed before showing up for lab. If the pre-lab assignment is not completed, you will not be permitted to perform the experiment until the pre-lab is complete. If as a result you are unable to complete the experiment before the end of the allotted time, you will need to make arrangements to complete the experiment AFTER SCHOOL within one week of the original date of the experiment or the experiment will be considered incomplete. The lab will be scored based on what is in the notebook at the end of that week. (WARNING: This after school lab makeup policy tends to make coaches and directors upset with students who are late to practice as a result. Keep this in mind should you decide to leave your pre-lab go until the last minute.)

1. Record the title of the experiment in the table of contents as well as the starting page number where it may be found.

2. On the first line of the experiment page, write and underline the experiment title.3. On the second line, include the date the experiment will be performed. If the experiment is

performed on multiple dates, add those dates here when they occur.4. Leave the third line blank for the names of your lab partner(s) to be filled in during the

experiment.5. Skip the fourth line and label the fifth line Purpose: and give a brief description of what the

experiment is and why it is relevant to what you are studying in chemistry. Do not copy the purpose from the experiment guide directly. Put the purpose into your own words to show that you understand the idea and to prevent plagiarism.

6. Skip a line, label the next line Hazards and Precautions: and list any special hazards you will encounter during the experiment such as corrosive and flammable chemicals as well as the chemicals that pose those hazards.

7. Skip a line, label the next line Procedure: and give a summary of the procedures to be used during the experiment. This should not be a detailed list of step by step procedures but should instead be a summary of what you will be doing. The laboratory manual itself will be used for the step by step procedures.

8. Skip a line, label the next line Pre-Lab Questions: On the next line, begin numbering and answering any and all prelaboratory questions at this time. It is not necessary to re-write the questions. Simply answer them using complete sentences when appropriate and showing all work whenever a calculation is required.

9. Skip a line after the last question and label the next line Results and Observations.

During the experiment1. Under Results and Observations: record all results and observations using as descriptive words

when possible. Example: “pale blue solid” or “sky blue powder” instead of simply “blue solid”2. Be certain to mention any time that something changes appearance or if bubbles, smoke, flame, or

other signs of chemical reactions occur.3. After completing the experiment, skip a line and label the next Data.4. In this section you will organize data into easy to view tables and charts as necessary. Be sure that

all measurements include the required units.5. After the data section is complete, skip a line and label the next Calculations:6. In this section you will show an example of all each kind of calculation performed during the

experiment. These calculations should include any formulas used and units on all measurements.

After the experiment1. Skip a line and label the next Conclusions:. Write a brief explanation of what can be concluded

based on the experiment. If you are unsure, look back to the Purpose section and see how the experiment accomplished that Purpose.

2. Skip a line and label the next Sources of Error: Write an explanation of specific sources of error in your experiment and how they affected your data. Did they cause your results to be larger or smaller than they should be? Human error should not be listed unless a specific mistake caused a significant result. If that is the case, explain why you did not go back and correct the error. Errors in calculations are also not a valid source of error. It is expected that if you know enough to know a calculation is wrong, you also know enough to either go back and fix it or to ask for help if you are unsure of how to fix it.

3. If it applies to the experiment, skip a line and label the next Percent Error:.4. Using the formula below, calculate the percent error, showing your calculations in your notebook.

Percent Error = |(Observed value – Accepted value)| x 100% Accepted value

5. After you finish the calculation, explain why you think your result has the degree of error that you calculated.

6. Skip a line and label the next Post-Lab Questions:.7. Beginning on the next line, number and answer the questions. It is not necessary to rewrite the

questions. If a question requires calculations rather than words, be sure to show all of your work.

Example notebook entry from Mr. Rickard’s notebook

Determining the Density of a Block of MetalFebruary 31, 2011Lab Partner: Ms. Cipolla

Purpose: The purpose of this experiment is to use experimentally determined mass and volume to calculate the density of a block of metal and then compare that to the known value for the metal. The formula used for density calculations will be density=mass/volume.

Hazards and Precautions: There are no chemical hazards for this experiment.

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Procedure: Ruler method – In the first part of the experiment, the metal sample will be massed on a balance and a ruler will be used to determine the dimensions of the metal so that volume may be calculated with the formula V = LxWxH.

Graduated cylinder method – In the second part of the experiment, the metal sample will be massed on a balance and it will then be submersed in a graduated cylinder to determine the volume of the metal by liquid displacement.

Pre-Laboratory Questions:1. Density is equal to the mass of a sample divided by its volume.2. The area of a box can be calculated by multiplying length times width times height.3. When an object sinks into the water and does not float, the water level will rise by an amount equal to the volume of the object.

Results and Observations: Metal sample is a copper cube Metal has shiny, orange color Mass = 2.34g Cube length = 0.83cm Cube width = 0.84cm Cube height = 0.85cm Starting volume of water = 5.90mL Ending volume of water = 6.44mL

Data:Method Mass Volume DensityRuler 2.34g 0.59cm3 3.9g/cm3

Cylinder 2.34g 0.54mL 4.3g/cm3

Calculations:

Ruler method volume V= LxWxHV = 0.83cm x 0.84cm x 0.85cm = 0.59cm3

Ruler method density D =M/VD = 2.34g / 0.59cm3 = 3.9g/cm3

Cylinder method volume V = final volume – initial volumeV = 6.44mL – 5.90mL = 0.54mL

Cylinder method density D = M/V

Conversion 1mL = 1cm3 so V = 0.54cm3

D = 2.34g / 0.54cm3 = 4.3g/cm3

Conclusions: In this experiment the density of copper was determined to be 3.9g/cm3 by the ruler method and 4.3 g/cm3 by the graduated cylinder method. Calculating the volumes in different ways did have an impact on the density even though the volumes seemed to be rather similar.

Sources of Error: The major sources of error in this experiment were the precisions of the balances, rulers, and graduated cylinders. Had more precise tools been available, this would have reduced the errors in the measurements.

Percent Error:

Percent Error = |(Observed value – Accepted value)| x 100% Accepted value

Accepted Value 8.9g/cm3

Ruler method = |(3.9 – 8.9)| / 8.9 x 100% = 56% error

Cylinder method = |(4.3 – 8.9)| / 8.9 x 100% = 52% error

Both results show a significant amount of error but are consistent with each other. This may indicate that there is an error with the calibration level of the balance. Multiple massings of the block may have shown an error in the original mass that was never double checked.

Post-Lab Questions:1. It should be possible to determine the identity of an unknown metal by calculating the density

of the unknown and then comparing the result to known values for various metals.

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Laboratory Procedures

How to use a Bunsen burnerBunsen burners are one of the most common sources of heat in high school chemistry laboratories.

They are either on or off and temperature control is accomplished by adjusting flame type or the object’s distance from the flame. Bunsen burners should not be used in a laboratory where volatile, flammable chemicals are being used.

Lighting Procedures1. Safety Check: Make sure the gas is off (handle perpendicular to the gas tap itself,) the tubing

is connected, goggles are on, long hair is tied back, and flammable objects are away from the Bunsen burner.

2. If the Bunsen burner has an air flow collar at the bottom with small holes, make sure that the collar is rotated so that the holes are covered/closed.

3. If the Bunsen burner has a valve that controls gas flow, tighten the valve the entire way and then turn it two or three turns to open it to gas flow. This is a starting point and it may be turned later to allow more or less gas to the burner.

4. Make sure you have matches or a striker ready to use.5. If you are using a match, light the match first then turn on the gas tap.6. If you are using a striker, check for usable flint, then turn the gas on, count to about five, and

then start striking. Remember to apply upward pressure on the striker to produce sparks.7. Adjust the flame by using the air flow collar or the control valve as needed.8. DO NOT leave a Bunsen burner unattended. If you need to leave your station, turn the burner

off. It can always be lit again later.9. Do not blow out a Bunsen burner flame. Turn the flame off by turning off the gas at the tap.

Flame TypesYellow/Orange Flame: This flame has the lowest temperature over around 300̊C and should not be

used for heating. It is primarily used to ensure the Bunsen burner is working.Large Blue Flame: This flame can be hard to see when the lights are on. It is very useful and is

about 500̊C.Inner Blue Cone: This is the hottest flame possible with a Bunsen burner reaching about 700̊C..

The blue flame is concentrated in a small cone near the top of the burner and can actually be heard. Making this blue cone flame takes practice and patients but is worth the effort.

How to use a hot plateHot plates are a useful alternative to Bunsen burners especially when multiple objects need to be

heated at the same time or when more precise temperature control is needed. WARNING: Due to their ceramic tops, hot plates stay hot for a LONG TIME after being turned off. Extreme care is required.

Heating Procedures1. Place the objects to be heated on the hot plate BEFORE the hot plate is turned on. This will help

to prevent glass from breaking due to heat stress caused by large temperature differences.2. Turn the hot plate on to the required setting.3. Monitor the hot plate while heating is occurring.4. DO NOT pour cold water on the ceramic top of the hot plate. The extreme difference in

temperature may cause the ceramic top to crack.

5. Turn the hot plate off when heating is complete. If possible, leave some indication that the hot plate had been used so that others do not accidentally burn themselves.

How to use filter paper for gravity filtrationOne of the easiest ways to separate small amounts of solids from solutions is by filtering. For

gravity filtration, you will need the following items:1. A funnel2. A piece of filter paper3. A ring stand and support ring to hold the funnel4. A container to collect the supernatant 5. The container with the original mixture6. A wash bottle with the same solvent as your mixture (usually this will be water)7. Patience

To perform a filtration using the above materials, the following steps may be used.1. Place the ring stand on the lab bench.2. Place the collection container on the ring stand base for stability.3. Attach the support ring to the stand such that the funnel stem extends into the collection

container.4. Move the collection container so that the funnel stem touches the inner wall of the

container to help minimize splatter.5. Fold the circular piece of filter paper in half.6. Fold the now half circle piece of filter paper in half again so that it looks like a quarter of a

circle.7. Along the curved edge of the folded filter paper, pull three edges to one side and leave the

fourth edge by itself. This will make a cone like cup to collect the solid without letting it pass through the paper.

8. Place the filter cone cup in the funnel and dampen with solvent from the wash bottle. This will help the filter paper stay in the funnel and will prime it for filtration.

9. Carefully pour the mixture into the filter cone to filter. Be careful not to allow the liquid level to rise above the top of the filter paper. This allows the mixture to bypass the filter paper and head directly to the collection container.

10. Wait patiently while the mixture filters, adding additional mixture to the funnel as it filters.11. Once filtering is complete, the liquid supernatant should be in the collection container and

the solids should remain in the filter paper.12. To allow the solids to dry more quickly, carefully remove the filter cone from the funnel

and unfold the cone to allow the solid to spread out and dry more quickly.

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How to take measurementsThe following procedures are meant to be used as quick reminders. If you do not understand how

to use a particular piece of equipment, check these procedures first with your lab partner. If you still cannot figure out the process, ask Your instructor for help.

MassMass is a measure of how much matter a substance has and is often reported with the units of

grams. We commonly INCORRECTLY refer to mass as weight. Try to break the habit of using weight when mass is the correct term. It will make you sound more intelligent and will also help you earn a higher grade when you take physics. If you would like a detailed explanation of the difference between mass and weight, sign up for physics next year.

BalanceZero Method

1. For liquids, place an appropriate container such as a beaker or flask on the balance to contain the liquid.

2. For solids, place a weighing paper or an appropriate container such as a beaker on the balance to contain the solid and protect the balance pan from the chemical.

3. Press the zero/tare button on the balance.4. Verify that the balance reads zero as appropriate.5. Add the chemical to be massed.6. Once the reading settles, record the mass value. The last digit of the mass has been rounded and is

therefore uncertain.

Container Mass Method1. For liquids, place an appropriate container such as a beaker or flask on the balance to contain the

liquid.2. For solids, place a weighing paper or an appropriate container such as a beaker on the balance to

contain the solid and protect the balance pan from the chemical.3. Record the mass of the container or paper. The last digit of the mass has been rounded and is

therefore uncertain.4. Add the chemical to be massed.5. Once the reading settles, record the combined mass value of the chemical in the container and the

container. The last digit of the mass has been rounded and is therefore uncertain. 6. The mass of the chemical is equal to the combined mass minus the mass of the empty container.

The last digit in the result is uncertain.

Example: Mass of Empty Beaker 121.21gMass of Beaker and Powder 123.45gMass of Powder 123.45g – 121.21g = 2.24g

VolumeVolume is a measure of how much space matter takes up and is commonly reported in liters or

milliliters. While each of the following may be used to measure volume, they each will vary in the amount of accuracy they provide based on the size, the number of graduation lines, and any special classifications of accuracy in the case of pipettes and burettes. Be sure to use the correct tool for the job.

Regardless of the container, the surface of liquids will have a slight downward curve called a meniscus. The correct way to determine the volume of liquid in a container is to make the measurement at the BOTTOM of the meniscus. Remember to estimate an additional decimal place past the smallest graduation line of the container. This estimated digit is uncertain.

Beaker - low accuracyBeakers are meant to hold chemicals rather than to accurately determine the volume of the liquid

in the beaker. Beaker graduations are meant as an estimate only.

Erlenmeyer Flask - low accuracyErlenmeyer flasks are meant to hold chemicals rather than to accurately determine the volume of

the liquid in the flask. Erlenmeyer flask graduations are meant as an estimate only.

Graduated Cylinder - medium accuracyGraduated cylinders are meant to provide and accurate estimate of the volume of liquid in the

cylinder before the liquid is transferred to another container. Graduated cylinders come in a variety of sizes and are more accurate than beakers and Erlenmeyer flasks but less accurate than pipettes, burettes, and volumetric flasks.

Pipette (aka Pipet) - high accuracyPipettes are used to transfer an accurate volume of liquid from one container to another. They use

a bulb to draw in the chemical. Never pipette by mouth. Pipets are more accurate than beakers, Erlenmeyer flasks, and graduated cylinders.

Burette (aka Buret) - high accuracyBurettes are used to transfer an accurate volume of liquid from one container to another at a

controlled rate of flow. They use a stopcock to control the rate of flow from a steady stream, to single drops, to no flow at all. Burets are more accurate than beakers, Erlenmeyer flasks, and graduated cylinders.

Volumetric Flask- high accuracy at single volumeVolumetric flasks are used to measure a single exact volume of a liquid or solution. While all of

the above glassware may have several graduations, a volumetric flask has a single reference line. The volume stated on the flask is the volume of liquid at that reference line. Accurate determination of volume above or below the reference line is not reliable. Volumetric flasks are more accurate than beakers, Erlenmeyer flasks, and graduated cylinders at the specified volume.

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TemperatureTemperature is a measurement of the amount of average heat energy of particles in a sample. In

the chemistry laboratory it is usually measured in degrees Celsius.

Digital Thermometer1. Place the thermometer in contact with the substance to be measured.2. Wait for the temperature readout to stabilize. The last digit has been rounded and is uncertain.

Liquid Thermometer1. Place the thermometer in contact with the substance to be measured.2. Wait for the liquid level inside of the thermometer to stabilize. The liquid in our thermometers

is alcohol and not mercury. If the thermometer breaks it is a fire hazard. 3. Read the temperature like you would the volume of a liquid. Remember to estimate an

additional decimal place past the smallest graduation line of the thermometer. This estimated digit is uncertain.

4. Be careful not to overheat the thermometer. If the volume of the liquid in the thermometer expands past the capacity of the thermometer, the thermometer will burst.

LengthLength is a measurement of how long a given object is in a given direction. In the chemistry

laboratory, length is typically measured in centimeters. Large distances may be measured in meters.

Ruler1. Make sure that the zero mark is lined up with one edge of the object to be measured. If the

ruler does not have a visible zero mark, make note of the measurement at the edge of the object or obtain a replacement ruler that does have a visible zero mark.

2. Using the graduation lines on the ruler, record the measurement as needed. Remember to estimate an additional decimal place past the smallest graduation line of the ruler. This estimated digit is uncertain.

Laboratory Experiments

This section of the laboratory manual will detail several of the major experiments that you will perform this year in chemistry. Your teacher may add additional experiments if time permits. Any additional experiments will follow a similar format and you will need to follow the normal policies and procedures.

The three golden rules for experiments are always1. Be safe in the lab.2. Known and understand the experiment before you do it.3. If you have questions, ask your instructor.

Golden Penny

Precipitates

Specific Heat

pH

Titration

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Experiment Title: Thinking Inside of a Box

Background Information:This first experiment is a simulated one where all of the data required has already been collected

and is provided to you. It is meant to give you a chance to learn how to write up an experiment in your laboratory notebook without the added stress of taking measurements and recording observations while performing the experiment yourself. Take your time. Follow the guidelines. Don’t wait until the last minute. If you have questions, email Ms. Cipolla or Mr. Rickard and they will get back to you as soon as their summer schedules permit.

The Story So Far: A sweaty-toothed mad man has trapped you and your entire chemistry class inside of the laboratory and has done devious and nefarious things to make the room completely air tight. He has agreed to let you out of the room if you can tell him how long before the oxygen in the room will run out and show data and calculations to support your answer. He reminds you that air is actually mostly nitrogen gas and is only 21% oxygen gas. He also gives you a hint that an average person uses 1.75 liters of oxygen gas per minute in order to breathe.

Purpose:To determine how long the oxygen gas supply will last in a sealed classroom full of students.

Materials:Quant Item Description1 meter stick1 calculator

Hazards:There are no chemical or equipment hazards for this experiment.

Pre-Laboratory Questions:1. How many liters are there in one cubic meter?2. How would you calculate the volume of a room if the room is estimated to be a rectangular box?

Procedures:1. Determine the length of the room as accurately as possible using a meter stick. Record this value

in your laboratory notebook. Do not forget the units!2. Determine the width of the room as accurately as possible using a meter stick. Record this value in

your laboratory notebook. Do not forget the units!3. Determine the height of the room as accurately as possible using a meter stick. Record this value

in your laboratory notebook. Do not forget the units!4. Clean up and put all of your laboratory equipment back where it belongs.

Simulated Data Table:Copy the following data table into your laboratory notebook in place of measuring the dimensions

yourself. These will be the values used for all of the required calculations.

Room Length 15.243mRoom Width 9.697mRoom Height 3.544m

Required Calculations: (REMEMBER TO SHOW ALL OF YOUR WORK)1. What is the volume of the room in cubic meters?2. What is the volume of the room in liters?3. How many liters of breathable oxygen are in the room?4. If your class has twenty four students and one teacher, how many minutes will the oxygen

last? Round to the nearest minute for these calculations.

Percent Error:There is no percent error for this experiment.

Post-Laboratory Question:1. Explain how measuring the room as a simple, empty box affects the amount of oxygen in the room

as compared to a situation in which you took into account for all of the other objects in the room that take up space such as desks, chairs, benches, etc.

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Experiment Title: Candle Observations

Background Information:In this experiment you will be exploring the differences between observing and seeing as well as

between observations and explanations. A wise chemistry teacher once said that the difference between seeing and observing is the same as the difference between hearing and listening. Listening and observing require active participation and attention by the person which hearing and seeing can be passive actions. By the end of the experiment you will hopefully find that by actively observing a candle, something you have probably seen dozens of times during your life, you will be able to develop predictions and theories about what is happening.

Purpose:1. To determine what observations can be made about how candles burn2. To determine if it is possible to predict how a candle will burn in a given situation3. To determine how to these observations and predictions differ from explanations of why the

candle burns in a certain way

Materials:To complete this experiment, you will need the following items

PER PERSON1 safety goggles 1 laboratory apron

PER GROUP1 pack of matches 1 ceramic dish1 600mL or larger beaker 1 400mL beaker1 ~100mL beaker 1 small crucible2 candles, tea light style 1 stopwatch or other timing

device

Hazards:You will be working with flames in this experiment. Therefore, all fire safety rules are in effect.

Be sure you know where all fire safety equipment is located before you start the experiment.Matches are not to be thrown in the sink or in the garbage can. They are to be collected in the

larger ceramic dish while they cool. At the end of the experiment, your teacher will explain what to do with the used matches.

Pre-Laboratory Questions:1. Where are the fire extinguishers, fire blanket, and safety shower located in your laboratory?2. What is the difference between an observation and an explanation?3. What three things are required for a fire to occur?4. What happens if one or more of the three things needed for a fire is missing or runs out while a

fire burns? 5. If you place two candles at different heights within a closed container, which one would you

expect to go out first? Explain your reasoning.6. ACADEMIC CHEMISTRY: Give a brief summary of the procedures of the experiment listed

below.

Procedures:

Part 1: Single Candle1. Place a single candle on the laboratory bench.2. Record at least three observations about the candle.3. Use a match to light the candle and then place the used match in the ceramic dish.

RECOMMENDATION: Do not blow the match out once the candle is lit. Instead, gently shake it back and forth until the flame goes out. This will prevent you from accidentally blowing out your candle or the candle of another group.

4. Record at least five observations of the burning candle while it burns for at least one minute. Do not move on to the next part until your group has made at least five observations.

Part 2: Single Candle with Beakers of Different Sizes

1. Cover your burning candle from Part 1 with the 600mL beaker and record how long it takes for the flame to go out.

2. Record at least three observations of what happened inside of the beaker once the candle was covered with the beaker.

3. Remove the 600mL beaker and light the candle with a new match.4. Cover the burning candle with a 400mL beaker and record how long it takes for the flame to go

out.5. Record any differences you observe between what happened when the 400mL beaker was placed

over the candle compared to when the 600mL beaker was placed over the candle.6. Remove the 400mL beaker and light the candle with a new match.7. Cover the burning candle with a 100mL beaker and record how long it takes for the flame to go

out.8. Record any differences you observe between what happened when the 100mL beaker was placed

over the candle compared to when the 600mL and 400mL beakers were placed over the candle.

Part 3: Single Candle in a Reused Beaker

1. Rinse the 400mL beaker out with water and then dry it with a paper towel.2. Light the candle with a new match.3. Cover the candle with the 400mL beaker and start the timer. 4. Remove the beaker carefully when the candle is ALMOST out and stop the timer.5. Sit the 400mL beaker FACE DOWN on the laboratory bench as gently as possible.6. Light the candle again if it accidentally went out in step twelve.7. Lift the 400mL beaker straight up and use it to cover the candle. Record how long it takes for the

candle to go out.

Part 4: Two Candles in the Same Beaker1. Turn the crucible upside down and place it on the laboratory bench.2. Place one of the candles on top of the crucible.3. Place the other candle on the laboratory bench beside and touching the crucible.4. Check to ensure that you can cover both candles at the same time with your large beaker. If you

cannot, double check your setup and get a larger beaker if necessary.5. Lit both candles with a new match.

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6. Cover both candles at the same time with the single large beaker and record how long each candle flame takes to go out.

7. Clean up your laboratory station and return all materials to the designated locations.

Example Data Tables:Beaker Size Time Burning 400mL Beaker Time Burning

First UseSecond Use

Candle Location Time BurningOn CrucibleOn Bench

Percent Error:There are no percent error calculations for this experiment. You will still need to discuss sources

of error.

Post-Laboratory Questions:1. What relationship did you observe between the size of the beaker and the amount of time it took

for the candle to go out?2. HONORS ONLY: Draw a graph of TIME vs BEAKER SIZE from Part 2 to show this

relationship.3. Develop a theory based on your observations from Parts 1 and 2 to explain WHY the candle goes

out.4. Using your theory from Post-Laboratory Question #3, explain your results to experiment Part 3.5. Compare your results from Part 4 to the prediction you made in Pre-Laboratory Question #5. If

the results were different than what you predicted, explain why the candles went out in the order that they did.

6. Often times on the inside of the beaker a hazy fog is observed. What is this fog and where does it come from?

7. HONORS ONLY: Remember, your answer to Post-Laboratory Question #6 is a theory. Develop an experiment to test the correctness of your theory.

Experiment Title: Separating the Sludge

Background Information:There are a variety of methods that may be used to separate components in a mixture. No one

method will work in all situations because of the variety of physical and chemical properties of the components. When liquid chemicals have distinct polarities like oil and water, they may separate into layers that can be separated by hand. Some solid components may be large enough that they may be removed by hand or with a pair of tongs or tweezers. Smaller components may be separated by using a filter much like a coffee pot. Sometimes the particles will be too small to filter or may be dissolved in another component and a process called distillation is needed. Distillation relies on the substances having different boiling points so that one boils away and leaves the other behind in the container.

Purpose:To separate a mixture into its components using processes such as filtration and distillation.



PER GROUP1 hot plate 1 laboratory balance2 250mL beaker 2 watch glass to cover beaker1 insulating pad or wire mesh 1 laboratory tongs1 filter paper 1 clay triangle1 ring clamp 1 glass stir rod1 ring stand 1 funnel1 tweezers 1 beaker tongs

Sludge Sample weighing paper

Hazards:This experiment makes use of hot plates, hot water, and hot glass. Proper procedures for the

prevention of burns, broken glass, and cuts must be observed.

Pre-Laboratory Questions:1. Explain how filtration works.2. Describe a situation in which filtration would not work and explain why filtration would not work.3. Explain how distillation works.4. Is it possible to collect all components of a mixture separated using distillation? Explain your

reasoning.

Procedures:

Part 1: Removing Solids1. Obtain a Sludge Sample from your instructor.

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2. Determine the mass of the Sludge Sample as accurately as possible and record the value.3. Record the mass of a clean piece of weighing paper.4. Remove any large chunks from the Sludge Sample and place them on the weighing paper. Make

sure there are no small particles stuck to the large chunks. Any small particles should be returned to the Sludge Sample.

5. Record the mass of the chunks and the weigh paper.6. Place the remaining Sludge Mixture in a 250mL beaker and add approximately 100mL of water.7. Stir the mixture thoroughly to make sure any particles that will dissolve have done so.8. Using a PENCIL, write your Period and Group Numbers on both sides of the

filter paper so that you will be able to identify it in the future.9. Record the mass of a piece of filter paper.10. Prepare the filter paper for gravity filtration inside of the funnel. See the

instructions in the Laboratory Procedures section earlier in this manual if you need a step by step procedure.

11. Record the mass of the second clean beaker.12. Construct the experimental setup pictured in Figure SS1 using a ring stand, a

clean beaker, ring clamp, clay triangle, funnel, and filter paper.13. Decant the liquid from the first beaker into the filter using a stirring rod. DO NOT allow the

liquid level to rise above the top of the filter paper. Try to keep as many of the solid particles as possible in the original beaker for now to help reduce the filtration time.

14. Move the solid particles to the filter paper and rinse the beaker with a small amount of water to transfer the last few particles if needed.

15. Set the beaker with the supernatant (the liquid that passed through the filter paper) aside for Part 2 of the experiment.

16. Carefully remove the filter paper from the funnel and sit it on the watch glass. 17. Open the filter paper so that the sand spreads out and dries as quickly as possible.18. Store the filter paper as instructed so that it will dry over night so that you may determine its mass

tomorrow.

Part 2: Recovering Dissolved Components1. Record the mass of a second clean, dry watch glass.2. Set up a hot plate on your bench.3. Place the beaker with supernatant you set aside in Part 1 on the hot plate and cover it with the

watch glass.4. Turn the hot plate on at a setting of between 8 and 10.5. Watch the beaker as the solution begins to boil. 6. Reduce the heat to a setting between 4 and 6 once the solution is boiling.7. Remove the beaker from the hot plate using the tongs when the water has evaporated. If the salt

crystals start spattering and popping, it is time to take the beaker off the hot plate. The watch glass is there to keep the salt inside of the beaker.

8. Allow the beaker to cool to room temperature.9. Record the mass of the beaker, watch glass, and salt.

Part 3: Clean Up1. Make sure that your wet sand and filter paper is located in the proper area for drying.2. Wash all glassware in the sink. The salt solution may be rinsed down the drain.3. Return all equipment to the proper location in the laboratory. If the hot plate is still hot, simply

leave it on your laboratory bench while it cools.4. Wipe down your bench using wet paper towels.

Figure SS1

5. Return your safety equipment to the appropriate locations.

Part 4: Mass of Sand and Final Clean Up1. Record the mass of the dry sand and the dry filter paper. If your filter paper is still damp, place it

in the appropriate location per your instructor and check it the following day.2. Place the dry sand in the appropriate container per your instructor and place the filter paper in the

garbage.3. Wash the watch glass and return it to the appropriate location.4. Wipe down your lab bench using damp paper towels.

Example Data Tables:

Experimental DataMass of Sludge SampleMass of weighing paperMass of weighing paper with chunksMass of dry filter paperMass of dry filter paper with dry sandMass of second clean beakerMass of second watch glassMass of beaker, watch glass, and salt

Required Calculations:1. Calculate the mass of the chunks.2. Calculate the mass of the sand.3. Calculate the mass of the salt.4. Calculate the total mass of all collected components.5. Calculate a percent mass of the chunks.6. Calculate a percent mass of the sand.7. Calculate a percent mass of the salt.

Percent mass formula: component mass x100%Sludge Sample mass

Percent Error:Calculate the percent error between the total masses of all components to the original mass of the

sludge mixture.Formula:

Sludge Sample mass – (collected component mass) x100%Sludge Sample mass

Post-Laboratory Questions:1. Why was it necessary to filter the mixture before the water was boiled?2. Why was it unnecessary to determine the mass of the water at any point during the experiment?3. Why was it important to use a watch glass to cover the boiling supernatant?4. Why was it necessary to allow the sand to dry over night before obtaining the mass?5. How would using a mass of damp sand have affected your results?

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Experiment Title: Measurement Making

Background Information:In the past, you have most likely used the English System of measurement which includes units

such as inches, pounds, and degrees Fahrenheit. Unfortunately, these are not the most commonly used units within the science laboratory. The Metric System is used for most of the standard measurements in science laboratories worldwide. This helps scientists to better understand results obtained by other scientists regardless of where the scientists live and work.

In math classes such as algebra and geometry, numbers can simply be numbers. For example, division problem answers will often be left in fractional form such as 4/3 or will use an indication for a number that repeats indefinitely such as 1.3. In the science laboratory however, most numbers we are using in our calculations are not just numbers, but are actually measurements. Measurements all have a level of precision and error based on the equipment being used to make the measurement. This often makes the fractional or repeated answer inaccurate. The key is to determine the correct level of precision and record the measurement to reflect that precision.

Purpose:1. To review the Metric System of measurement2. To practice measuring using common laboratory equipment3. To practice determining error levels of measurements.



PER CLASSVarious stations around the room will be used. These will be provided by the instructor. Do not alter the stations at any time.



Pre-Laboratory Questions:1. Review the How to Take Measurements section at the beginning of this manual.2. Explain what a meniscus is and why it is important.3. Explain why estimation is necessary for many measurements in the laboratory.

Procedures:

For this experiment, you will start at one of several stations located around the room, each with its own set of instructions. You will have a limited time at each station before rotating to the next station. Be sure to make all of the indicated measurements at each station. This experiment will be due at the end of the period and may not be taken home without permission from the instructor.

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Mass Station #11. Record individual masses for each of the three items at the station.2. Record a combined mass of all three items by placing them on a balance at the same time.3. Explain why it is possible for the sum of the individual masses not to equal the mass obtained by

massing all three items at the same time due to the accuracy of the balance itself.

Mass Station #21. Record individual masses for each of the three items at the station.2. Determine what combination of items would provide a total between 20.0g and 21.0g. (Example,

if Item#1 is 3.0g, Item #2 is 2.0g, Item #3 is 25.1g, one possible solution would be to use seven of Item #1 for a total of 21.0g)

3. Determine a second combination that would meet the same requirements.

Volume Station #11. Record individual volumes for each of the three liquid samples located in graduated cylinders at

the station.2. Add 200mL of water directly to the 600mL beaker at the bench.3. Transfer the water from the beaker into a 250mL graduated cylinder and record the volume as

measured in the graduated cylinder.4. Repeat steps 2 and 3 two more times.5. Explain whether the beaker or the graduated cylinder is more accurate and be sure to explain your

reasoning.

Volume Station #21. Record individual volumes for each of the three liquid samples located in graduated cylinders at

the station.2. Measure out one cup of water as exactly as you can using the kitchen style measuring cup

provided.3. Transfer the water from the measuring cup into a 250mL or 500mL graduated cylinder and record

the volume as measured in the graduated cylinder.4. Repeat steps 2 and 3 two more times.5. Calculate an average volume in mL for one cup of water.6. Show your answer to your instructor who will give you the accepted value of mL in one cup.

Calculate the percent error between your average and the accepted.

Temperature Station 1. Record the temperature of the boiling water using each of the three thermometers at the station.2. Record the temperature of the ice water using each of the three thermometers at the station.3. Explain which thermometer you think is the most accurate and how you reached that conclusion.

Length Station #11. Record the length of each of the three thin, straight items located at the station.2. Record the circumference of the beaker at the lab station.3. Calculate the area of the piece of aluminum foil at the lab station.

Length Station #21. Record the dimensions of a floor tile using each ruler located at the station.

2. Estimate the area of the classroom.3. Explain how you arrived at your estimate.

Combined Station1. Determine the volume of the metal cube using the ruler provided at the station.2. Determine the volume of the metal cylinder using the ruler provided at the station.3. Determine the volume of the metal cube using the graduated cylinder provided at the station.4. Determine the volume of the metal cylinder using the graduated cylinder provided at the station.5. Explain whether the ruler or the graduated cylinder is more effective for measuring volumes of

small objects and how you reached that conclusion.


Station Object Description Measurement

Required Calculations:All required calculations were listed in the Procedure section of the experiment.

Percent Error:Show your percent error calculations for the measuring cup of water from Volume Station #2.

Post-Laboratory Questions:All required questions were listed in the Procedure section of the experiment.

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Experiment Title: Density

Background Information:Objects can come in various shapes and sizes. There could be a large wooden block or a small

rubber ball. While the size and shape of two objects may be different, if they are made of the same substance, the two objects should have the same density because density is a physical property of a substance. Density is a measure of how much mass a sample has divided by the volume of the sample. For any given substance, its density should be constant regardless of the size and shape.

This can be very helpful in identifying an unknown as you are about to experience for yourself during the course of this experiment..

Purpose:1. To experimentally determine the density of several metal cylinders.2. To determine the identity of an unknown metal based on its experimentally determined

density.



PER GROUP1 metal sample set 1 laboratory balance1 ruler 1 unknown metal sample1 10 or 25mL graduated

cylinder

Hazards:There are no chemical hazards in this experiment. Don’t lick the glassware. Don’t chew on the

metal cylinders or unknowns. Use common sense and you’ll be fine.

Pre-Laboratory Questions:1. Which would you expect to have a higher density golf balls or Styrofoam? Explain your

reasoning.2. How can you tell the difference between pennies, dimes, nickels, and quarters in a pile?3. What are two different ways of determining the volume of a small cube in a chemistry laboratory?

Procedures:

For each metal sample complete each of the following steps.1. Write down observations about color, shape, and other observable properties of the sample.2. Determine the mass of the sample on a balance and record it.3. Estimate the volume of the sample using a ruler. Record all measurements and show the

calculations.

4. Fill a graduated cylinder about halfway with water.5. Record the volume of water.6. Place the sample in the graduated cylinder carefully by tilting the cylinder and sliding the sample

gently down the side.

Clean-up1. Dry all metal samples, known and unknown with a paper towel.2. Return all glassware and supplies to the appropriate locations.3. Return all other supplies to the appropriate locations.4. Wipe down the laboratory bench.5. Return safety equipment to the appropriate locations.


Sample Mass Ruler Volume

Ruler Based Density

Graduated Cylinder Volume

Graduated Cylinder Based Density

Required Calculations:1. Calculate the volume of the sample based on the water volumes and record it.2. Calculate the density of the sample using both the ruler volume and the graduated cylinder volume

and record them.

Percent Error:1. Calculate the percent error between the densities that used the graduated cylinder volumes and the

accepted values for metals listed in class.2. Calculate the percent error between the density of your unknown and the three closest densities

from the chart of accepted values listed in class.

Post-Laboratory Questions:1. What substance do you think is your unknown? Explain your reasoning. Be sure to mention your

percent error calculation results as a part of your explanation. HINT: If you didn’t pick the smallest percent error as the identity of your unknown, you need to explain why you picked another metal.

2. Which method gave better results, the ruler volumes or the graduated cylinder volumes? Explain your reasoning.

3. How could traders tell the difference between real gold and “fool’s gold” using skills you have used in the laboratory?

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Experiment Title: Hydrates

Background Information:Have you ever ended up with a wet cell phone? Maybe you got stuck in the rain. Maybe you were

pushed into a swimming boots. Maybe you dropped it in the toilet while texting. Before taking it back to the store and hoping they will fix it for you, one of the things you can try is to put the damp phone in a bag with rice for a day or two. Over time, the rice will absorb moisture from the phone and hopefully the phone will work. An alternative to rice would be the silica packets that come with a lot of electronics. You know the one, it has the “do not eat me” warning on it. It is there to keep the electronics for getting damaged by moisture between the factory and when you buy it.

A lot of chemicals are like this. We call them “hygroscopic” meaning, they absorb moisture from their surroundings, including the humidity from the air. In the chemistry lab, sodium hydroxide is the classic example. You can see just how much moisture it absorbs by leaving it on a balance and watching it over time. That, however, is another lab for another day.

This lab will focus on chemicals that have extra water trapped in with the actual chemical. The chemical and water molecules come as a set. Depending on the chemical, it may come with a single water molecule in a 1:1 ratio. Other chemicals come in a 2:1 ratio. Still other chemicals may come with more than ten water molecules for each one molecule of chemical. That’s a lot of water! And that water adds mass.

Purpose:1. To determine the percentage of water in a chemical with known composition.2. To determine the percentage of water in a chemical with unknown composition.3. To determine the hydrated formula of the chemical with unknown composition.



PER GROUP1 hot plate 1 laboratory balance2 400mL beaker 1 watch glass to cover beaker1 insulating pad or wire mesh 1 laboratory tongs

calcium chloride dihydrate unknown sample

Hazards:calcium chloride dihydrate

Pre-Laboratory Questions:1. Write the formula for calcium chloride.2. Calculate a molar mass of calcium chloride.3. Explain why it is necessary to know whether or not a chemical is a hydrate if you are preparing a

2.0M solution.

Procedures:

Part A: Percentage of Water in Calcium Chloride Dihydrate1. Record the mass of the clean, dry beaker and watch glass.2. Remove the watch glass using the tongs.3. Zero the balance.4. Add calcium chloride dehydrate until the balance reads between 1.0 and 2.0 grams and then record

the mass.5. Move the beaker to the hot plate using the tongs.6. Cover the beaker with the watch glass using the tongs.7. Heat the hydrate gently to avoid spattering the salt until all moisture has evaporated and can no

longer be seen on the beaker or watch glass. The salt should change from crystals to a powder.8. Use tongs to transfer the hot beaker and watch glass to the insulating pad or wire mesh to allow it

to cool for at least five minutes. Do not touch it during the five minutes.9. Use tongs to transfer the beaker and watch glass to the balance. 10. Record the final mass of the combined beaker, watch glass, and powder.11. Empty the contents of the beaker into a paper towel and discard in the designated container. 12. Clean the beaker out with a dry paper towel. Do not use water at this point.13. Calculate the percentage of water in the sample based on your results and record on the

instructor’s sheet.14. Check with the instructor to see if you need to repeat Part A or if you may move on to Part B.

Part B: Percentage of Water in an Unknown Sample.1. Obtain an unknown from your instructor.2. Record the letter or number of your unknown.3. Record the mass of the clean, dry beaker and watch glass.4. Remove the watch glass using the tongs.5. Zero the balance.6. Add calcium chloride dehydrate until the balance reads between 1.0 and 2.0 grams and then record

the mass.7. Move the beaker to the hot plate using the tongs.8. Cover the beaker with the watch glass using the tongs.9. Heat the hydrate gently to avoid spattering the salt until all moisture has evaporated and can no

longer be seen on the beaker or watch glass. The salt should change from crystals to a powder.10. Use tongs to transfer the hot beaker and watch glass to the insulating pad or wire mesh to allow it

to cool for at least five minutes. Do not touch it during the five minutes.11. Use tongs to transfer the beaker and watch glass to the balance. 12. Record the final mass of the combined beaker, watch glass, and powder.13. Empty the contents of the beaker into a paper towel and discard in the designated container. 14. Clean the beaker out with a dry paper towel. Do not use water at this point.15. Calculate the percentage of water in the sample based on your results and record on the

instructor’s sheet.

Part C: Clean Up6. Wash all glassware with water and return to the appropriate locations.7. Clean off the pan of the electric balance.8. Clean the top of the hot plate once it has cooled. Do not try to wash the hot plate before it has

cooled so that the ceramic top does not shatter.9. Wipe down your lab bench with a damp sponge or paper towel.

44


Part A Part BMass of beaker and watch glass

Mass of beaker and watch glass

Mass of calcium chloride dihydrate

Mass of unknown hydrate

Combined mass of beaker, watch glass, and calcium chloride dihydrate

Combined mass of beaker, watch glass, and unknown hydrate

Combined mass of beaker, watch glass, and calcium chloride anhydrate

Combined mass of beaker, watch glass, and unknown anhydrate

Mass water lost Mass water lost

Required Calculations:1. Mass water lost for each trial.2. Percent water in sample for each trial.

Formula: mass of water x 100=mass of hydrate

3. Percent of water in CaCl2*2H2O

Percent Error:Calculate the percent error in your measurement of water in the calcium chloride dihydrate. The

accepted value for calcium chloride dihydrate is your answer to Required Calculation #3.

Post-Laboratory Questions:1. Why was it necessary to handle the beaker and watch glass with tongs after recording their

masses?2. How could the procedure be changed in order to ensure that all of the water has left the original

sample making it completely anhydrous?3. Based on your results from Part B, determine the identity of your unknown by comparing it to a

list provided by your instructor. Which chemical do you think was your unknown? Explain your reasoning.

4. HONORS ONLY: Determine the chemical formula of the hydrated version of your unknown. Show all of your calculations. Your answer should be in the form AxBy*ZH2O.

Experiment Title: Precipitation

Background Information:Many times people think that just because a solution is clear, it must be clean and not have any

chemicals in it. Nothing could be further from the truth. Many clear solutions have ions in them just waiting to react with other substances and ions. In this experiment, various solutions of various colors will be mixed in small amounts and reactions may or may not take place. Keep a sharp eye out and use the solubility rules to help determine what is reactions are occurring and which combinations are simply “No Reactions.”

Purpose:1. Observe and identify precipitates formed when aqueous solutions are mixed.2. Write balanced equations for each reaction.3. Determine which of the possible products is the precipitate based on solubility rules.



PER GROUP1 96-well plate for reactions 1 24 well plate for chemicals1 set of 10 chemicals in

disposable pipets1 sheet white paper

Hazards:sodium hydroxide:

barium hydroxide:

Pre-Laboratory Questions:1. What is a precipitate?2. How can you observe if a precipitate has formed?3. Use solubility rules to determine which of the following four chemicals would form an aqueous

solution and not be a precipitate. There may be more than one. List the rule number that led to your decision for each chemical.

KCl AgCl Al(OH)3 Ca(OH)2

Procedures:1. Obtain a set of sample chemicals from your instructor.2. Place your 96 well plate on the piece of white paper.3. Add the first sodium containing chemical to the first five wells of the first row.4. Add the second sodium containing chemical to the first five wells of the second row.5. Repeat with each sodium containing compound.

46

6. Add the first chloride containing chemical to each of the wells in the first column that already has a sodium compound in it.

7. Add the second chloride containing chemical to each of the wells in the second column that already has a sodium compound in it.

8. Repeat with each chloride containing compound.9. Create a data table like the one listed below.10. Record your observations of what happened in each well in the appropriate block of the data table.

Be sure to look at the well plate over the white paper and over the black laboratory bench to make sure you don’t miss anything.

Clean-up Procedures11. Place a paper towel on the laboratory bench.12. Turn your 96 well tray over onto the paper towel and then gently tap the tray against the bench to

dislodge the solutions in the wells.13. Throw the paper towel in the garbage can.14. Rinse the 96 well tray three times with tap water, tapping the tray against the inside of the sink

after each rinse to dislodge the water in the wells.15. Place your 96 well tray upside down on a clean paper towel to dry.16. Return all other supplies to the appropriate locations.17. Wipe down the laboratory bench.18. Return safety equipment to the appropriate locations.

Example Data Tables:CoCl2 CuCl2 AlCl3 BaCl2 NiCl2

NaI cloudy ppt, light blue

No RXN

Na2CO3

Na3PO4

Na2SO4

NaOH

Post-Laboratory Questions:1. Write a balanced chemical equation for each well in which a precipitate formed.2. Write a balanced complete ionic equation for each well in which a precipitate formed.3. Write a balanced net ionic equation for each well in which a precipitate formed.4. For each precipitate, list which solubility rule shows that the precipitate should be insoluble.

Experiment Title: Specific Heat

Background Information:This experiment will focus on the flow of heat energy between objects of temperatures until

equilibrium is established. It relies on the property of specific heat being consistent for a substance regardless of the size of the sample of the substance.

Purpose:1. Experimentally determine the specific heat values for metal samples of known composition

using calorimetry2. Calculate the accuracy of the experimental results by comparing to accepted values.

Materials:Appropriate safety attire is required.

Quant Item Description Quant Item Description1 coffee cup calorimeter 1 digital thermometer1 400mL beaker for tap water

sample1 electronic balance, higher

precision preferred1 25mL graduated cylinder 1 metal samples

Hazards:Hot plates and hot water will be used in this experiment.

Pre-Laboratory Questions:1. none at this time

Procedures:1. Fill a beaker with 300mL of tap water and sit it aside so that it comes to room temperature.2. Set up a coffee cup calorimeter3. Record the mass of your calorimeter.4. Measure out about 75mL of tap water and add it to the calorimeter.5. Record the mass of the combined calorimeter and water.

THE FOLLOWING STEPS NEED TO BE CARRIED OUT QUICKLY AND EFFICIENTLY6. Record the temperature of the water and calorimeter.7. Obtain a hot metal sample from one of the hot plates.8. Add the metal sample to your calorimeter and stir until the temperature stabilizes.9. Record this temperature in your laboratory notebook.10. Remove the metal sample from the calorimeter and dry it off completely with a paper towel.11. Record the mass of the metal sample.12. Empty your calorimeter water down the drain.13. Repeat steps 4-12 as instructed.

48

Required calculations: (For this lab, show all calculations for all trials.)

1. Calculate the energy gained by the calorimeter in each trial. The specific heat of water is 4.184J/gK and the specific heat for your calorimeter is as given by your instructor. (This will have been determined in a previous lab.)

2. Using Conservation of Energy, calculate the specific heat value for each of your metal samples.

3. Calculate your percent error between your experimental values and the values given to you by your instructor.

Percent Error:Percent Error = |(accepted value – your value) / accepted value| x 100%

Post-Laboratory Questions:1. Explain how the Law of Conservation of Energy makes this lab possible.2. Explain how the sign of “q” indicates whether heat was lost or gained.3. What could you have done differently to make this lab work more effectively? 4. What sources of error contributed to your values being different than the accepted?

REMEMBER: Human Error and Math Mistakes are not acceptable answers.

Conclusion:Write a one paragraph conclusion that covers the following key points:

your actual resultsyour percent errorshow accurate you think the results are within the confines of the experiment

Experiment Title: Calorimetry – Heat of Fusion of Ice

Background Information:Insert background information here

Purpose:To determine the amount of heat energy required to melt one gram of ice. This value is known as the heat of fusion and is measured in joules.



PER GROUP1 250mL beaker 1 scoop2 Styrofoam cups 1 100mL graduated cylinder1 thermometer 1 hot plate1 beaker tongs 1 support stand1 extension clamp sample of ice



Pre-Laboratory Questions:1. Explain what happens at a molecular level when ice melts.2. Explain what happens at a molecular level when water boils.3. What phase does water typically take at room temperature?4. How does sitting near a camp fire keep you warm?

50

Procedures:1. Place a beaker containing about 120mL of water on the hot plate.2. Support the thermometer in the beaker per Figure HFI1 so that you

may measure the temperature of the water as it is heated. Be sure not to have the thermometer touching the bottom of the beaker as you may get an incorrect temperature reading of the beaker rather than the water inside of the beaker.

3. Heat the water to a temperature of about 50̊C.4. Measure 100.0mL of the hot water using the graduated cylinder.5. Record this volume as V1.6. Pour the hot water into a Styrofoam cup.7. Measure and record the temperature of the water in the Styrofoam cup

as T1.8. Immediately add several ice cubes to the cup.9. Continue to add ice so that ice is present at all times.10. When temperature no longer decreases, this is T2. Record11. Carefully remove the remaining ice. Try not to take water with it.12. Find the volume of water in remaining in the cup.13. This volume is V2. Record.


V1

V2

T1

T2

Required Calculations:1. Find the volume of water that resulted from the ice melting. (V2-V1)2. Using the density of water -1.00g/mL, find the mass of 100.0mL of water and the mass of the

volume in #1.3. Find the change in temperature.4. Calculate the heat transferred using. ΔQ = mcΔT

c = 4.2J/g ̊Cm = answer found in #2ΔT = the answer found in #3

5. The heat of fusion of ice is determined by dividing the amount of heat transferred by the mass of water resulting from the ice melting.

Percent Error:Calculate your percent error for the heat of fusion of ice.

Accepted value is 336J/g

Post-Laboratory Questions:1. How much energy would be needed to melt 75.0g of ice based on your results? Show your

calculations.

Figure HFI1

2. How would having excess water on your ice at the start affected your results? Explain your reasoning.

3. How would having ice below 0 degrees Celsius affected your results?Explain your reasoning.

4. Why did you need to be careful to take as little water as possible out when you removed the ice?

52

Experiment Title: Molar Volume of Hydrogen Gas

Background Information:One mole of an ideal gas should occupy a volume of 22.4L according to the Ideal Gas Law at

STP. The challenge in experimentally testing gases in the real world becomes complex because the amount of gas measured must be experimentally produced rather than simply obtained from a gas cylinder or other traditional method of containment. The relatively light mass of a mole of most common gases adds an additional challenge. Due to the limitations of most laboratory balances, the mass of the gas produced, and thereby the moles, must be calculated rather than directly measured.

Purpose:1. To experimentally create a known quantity of hydrogen gas2. To determine the volume of the produced hydrogen gas3. To compare the created volume with the accepted volume of 22.4L per mole gas



PER GROUP1 600mL or larger beaker 25mL 1.0M hydrochloric acid1 50mL or 100mL burette 1 thermometer3.5cm magnesium ribbon tap water as needed1 burette stand 1 burette clamp1 rubber stopper for the burette 1 thin stem funnel

Hazards:Determine the hazards for the following chemicals:magnesium ribbonhydrochloric acidhydrogen gas

Pre-Laboratory Questions:1. Using the internet or a catalog, determine the cost of a 50mL burette with a Teflon or PTFE

stopcock. Your answer should include the company, price, and part number. If you break one, this will be handy information to have.

2. Using gas laws, explain why it is important to know what the temperature of the laboratory is when you are performing the experiment.

3. Write a balanced equation for Mg(s) reacting with HCl(aq) to form aqueous magnesium chloride and hydrogen gas.

4. ACADEMIC ONLY: Give a brief summary of the experiment listed below.

Procedures:1. Setup the stand, clamp, beaker and water as per Figure MHG12. Obtain a piece of magnesium ribbon that is approximately 3.5cm

longFigure MHG1

3. Determine the mass of the piece of magnesium using the balance with the most number of decimal places available. It should be somewhere between 0.02g and 0.04g. If not, alter the size of your strip or obtain a new strip if needed.

4. Make sure the burette stopcock is closed and add about 25mL of the 1.0M HCl to the burette.5. Add 25mL of water to the burette carefully to prevent unnecessary mixing of the chemicals.6. Bend the magnesium ribbon into a U shape and insert it into the

burette as per Figure MHG2.7. Cap the burette with the rubber stopper and quickly put the burette

in the beaker end of the water, stopper end first. The stopper must be below the level of the water.8. Remove the stopper from the burette and check the liquid level in the burette. If the water level is

in the graduated region of the burette, record this starting volume. If the water level is still above the graduated region, open the stopcock just long enough to lower the water level until it is in the graduated region. Record this starting volume.

9. Gently lower the rim of the burette to the bottom of the beaker in case the magnesium ribbon falls to the bottom.

10. Observe and record what happens inside the burette during the reaction.11. Record the final volume in the burette once the reaction has reacted completion.


mass Mg ribbonstarting volume final volumevolume H2(g) producedmoles Mg used

Required Calculations:1. Calculate the volume of H2(g) produced.2. Convert the volume of H2(g) produced from milliliters to liters.3. Convert the mass of Mg to moles of Mg.4. Divide the volume of H2(g) produced by the moles of Mg to get the volume per mole ratio.5. Use Charles’s Law to determine an adjusted volume of H2(g) produced.6. Divide the adjusted volume of H2(g) produced by the moles of Mg to get the adjusted volume per

mole ratio.

Percent Error:Calculate the percent error for both ratios you calculated. The accepted volume per mole ratio is 22.40L/mol.

Post-Laboratory Questions:1. Why was it theoretically necessary to adjust the volume of gas produced for temperature?2. Did the temperature adjustment make a significant change in the percent error? Explain your

reasoning.3. Explain how having an electronic balance that reads to an additional decimal place should result in

a lower percent error for the experiment.

Figure MHG2

54

Experiment Title: pH


Purpose:Insert purpose here.

Materials:Insert materials here, chart preferredQuant Item Description Quant Item Description

Hazards:Insert hazard warnings and/or list hazardous chemicals here for students to look up for themselves.

Pre-Laboratory Questions:5. Insert pre-laboratory questions here.6. And here7. And here8. Etc.

Procedures:14. Insert procedure list here15. And here16. Etc17. And so on

Example Data Tables:Insert example data tables if needed

Percent Error:List all results requiring percent error calculations here

Post-Laboratory Questions:4. Insert post-laboratory questions here5. And here6. Yada yada

Experiment Title: Candy Coating Chemistry

Background Information:Almost any child will tell you that yellow and blue make green. What they mean is that when the

two primary pigment colors yellow and blue combine they result in a secondary pigment color typically described as green. In addition, when yellow is combined with the third primary pigment color, red, the result is orange. Finally, when blue and red are combined, violet results. Interestingly, when any of the secondary pigments are combined with the remaining primary pigment color, it results in a brown, a fact that can easily be confirmed by watching children using water color paints and a clear plastic cup to rinse their brushes.

When the three primary colors are combined in various ratios, the results span the spectrum from fire engine red, to a yellow orange, to teal, and a deep purple. Any easy place to see the variety of possible colors is in the candy industry. Candy is given a variety of colors in addition to shapes and sizes in order to make the candy stand out and seem appealing to potential consumers. The rules of pigment blending apply here as well.

The candy industry primarily relies on a set of standard food dyes and lakes which have been found to be safe for consumption. The major difference between dyes and lakes is that dyes are soluble in water and lakes are not. Dyes will provide color by “staining” an object or by simply being dissolved in the water portion of a product such as juice or soda. Lakes will provide color by dispersing themselves in fats and oils making them ideal for use in lipstick, soaps, cake mixes, and chewing gum. They also work well for hard coatings such as those found on prescription medication and M&M’s. By combining these various dyes and lakes in the correct combinations, the rainbow of colors one associates with a candy store can be created.

Chromatography is a process where mixtures are separated into their components. Most chromatography makes use of a stationary phase, which doesn’t move, and a mobile phase, which does. The mobile phase carries the components of the mixture through the stationary phase at different rates giving them a chance to separate. This separation allows components to be identified using a variety of methods including simple observation all the way up to complicated instrumentation that blows the components to bits and looks at the pieces.

In paper chromatography, the paper itself is the stationary phase. The mobile phase may be any liquid as long as it does not react with the paper. In this experiment, our mobile phase will be water with a bit of dissolved salt in it. As the water makes its way up the paper, it will drag the dyes with it to varying degrees. The distance the each dye travels compared to the distance that the mobile phase (water) travels is known as the Rf and should be a constant for each individual dye. The formula for Rf is:

Rf = (distance dye component) / (distance mobile phase)

Purpose:1. To experimentally separate the dyes in several common candy coatings via liquid chromatography 2. To observe the differences between dyes and lakes3. To determine the which colors of candy are safe for people who have an allergy to yellow food

dye

Materials:PER PERSON

1 safety goggles 1 laboratory apron

PER GROUP

56

2 600mL or larger beaker 4 watch glasses1 ruler 4 pieces of filter paper1 pencil1 scissors tap water as needed1 stapler candy samples1 small beaker paper towel for blotting1 glass rod pinch of salt

Hazards:While none of the candy samples are inherently dangerous, since they are being used in the

laboratory setting, none of the candy may be consumed at the end of the experiment.

Pre-Laboratory Questions:1. Using methods you have already learned about this year as well as the background information

above, how would you experimentally separate dye colorings from lake colorings?2. What is the importance of the mobile phase?3. What is the importance of Rf values?4. After you read the procedures, explain what you think will be the most difficult step and why you

think it will be difficult.

Procedures:1. Obtain all of the required materials from your instructor.2. Cut the edges off of each circle of filter paper in order to have four rectangles that are

approximately 9.0cm by 7.5cm.3. Using a pencil, draw a line 2.5cm away from one of the 9.0cm edges of each paper. This is the

line on which you will “spot” each of your samples so that they have an even starting point.4. Put a small amount of water in each watch glass.5. Dip the edge of one of the candies in the water in a watch glass.6. Touch the candy to a paper towel to remove excess moisture. The candy will still be damp.7. Touch the candy to the first filter paper sheets along the pencil line as neatly as possible. This

should leave a dot of the dye. The smaller and darker the dot, the better. 8. Touch the candy to the paper in the exact same spot a few times to get a dark dot.9. Place the candy in the water in the watch glass and leave it sit for a minute or two. Then flip it

over and leave it sit for another minute or two.10. Repeat steps 5 through 9 for each color of the candy until you have four spots on the first piece of

filter paper. Be sure to use a different watch glass for each piece of candy.11. Once you have all four candies spotted on the same filter paper, curl the paper into a tube with the

line on the outside. 12. Staple the top and bottom of the paper without having the edges overlap as demonstrated by the

instructor.13. Add about 1.0cm of water to the bottom of the beaker.14. Add a pinch of salt to the water and stir until the salt is dissolved.15. Place the filter paper tube into the water with the pencil line at the bottom. Be careful not to tip

the filter paper over. It must remain standing as the water climbs toward the top.16. While you wait for the water to rise to about 1.0cm from the top of the tube, make your

observations of the dyes/lakes in each of the watch glasses.17. After you have made your observations, take your watch glasses to your instructor.

18. After your instructor gives you back your watch glasses, repeat the above steps for the second set of candy, a second filter paper, and the other 600mL beaker.

19. Once the water level has risen to within 1.0cm of the top of the tube, carefully remove it from the beaker, open the tube up, and lay it flat on your lab bench.

20. Carefully trace a line along the water line at the top of the tube.21. Measure the distance from the original pencil line to the top line.22. Record this value in your notebook.23. Mark a small line beside the center of each die smear on the filter paper. You may have more

than one color smear for each piece of candy. If so, label the center of each color separately as best you can.

24. Measure the distance to each mark from the original pencil line and record those values in your lab notebook.

25. Repeat the above steps for each candy sample.26. Record observations of the instructor’s station as per instructor.

Example Data Tables:Filter paper #1Candy color Dot color DistanceBlue blue 3.5cmOrange Red 1.5cm

Yellow 0.5cm

Required Calculations:1. Calculate an Rf for each dot on each filter paper. Show your work.2. Create a chart that shows which dyes are present in each candy sample.3. Include a column in the chart that labels candy containing yellow dye as an allergen.

Percent Error:There are no percent error calculations for this experiment.

Post-Laboratory Questions:1. Why did the line need to be drawn using a pencil and not a pen?2. Why did you need to dip each candy piece into a separate watch glass?3. How would the tube falling over, even momentarily, have affected the results?4. Which candies contained a yellow dye and could therefore be potentially dangerous?5. Explain how the concept of “like dissolves like” affects the Rf value of a spot in paper

chromatography. ` `

Conclusion:For this lab, make sure your conclusion contains a summary of what you did, what results you

obtained and how accurate and reliable you think the results are.

58

Experiment Title: Strong Acid and Strong Base Titration




Hazards:Insert hazard warnings and/or list hazardous chemicals here for students to look up for themselves.






The Golden FleeceIf it is gold that you want do your works and leave town before sunrise.

Bathe the lowliest coin of the realm in acid muriaticus,till it shines with the luster of the metal cupras.Wash in pure spring water, let stand till needed.

Place a small spoonful of the metal the Prussians call zink, in a glass vessel. Cover with aqua alkali, which is soda lye, and heat till steam begins.

Plunge the coin into the steaming bath.Flammable air begins to rise and the token will turn to purest silver before your very eyes.

Washed and dried, the silver token is now held in the outer flame of a strong fire. In a blink the change to gold will fill your eyes with wonder.

Blink twice more, then plunge the token into a bath of spring water. You now have gold, fit for the finest works.

^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^ Translate the procedure above following these steps:

1. Write the entry for this lab in your lab notebook following the format prescribed by your instructor. (Title, purpose, materials, precautions, procedure, observations, conclusion/post-lab questions)

a. Replace the ancient names of chemicals with their current scientific names.b. Use MSDS to write precautions and dangers for the chemical you will use.c. Write a step by step procedure using laboratory equipment from your lab drawers,

hot plates, and lab burners. Include handling equipment and safety equipment too.

d. Highlight precautions.e. Submit your notebook to your instructor for approval.

2. Complete the lab in the time allotted. You must make at least one gold coin to be taped in your lab notebook. 3. Write your conclusion for this lab after our class discussion.

^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^

60

White Powders Lab

Purpose: To perform a series of tests to determine the physical and chemical characteristics of several unknown powders. To utilize test results to identify an unknown substance.

Background Information:

Chemical High School has a drug problem. Over the past year, illegal drugs have been seized from student lockers on five occasions. All of these illegal drugs are white powders that look like table salt. During a recent locker search, investigators collected several vials filled with white powder. Before charges can be pressed on the individual in possession, the identity of the powder must be established.

You are a member of a forensic science lab team that has been sent to Chem High. The unknown white powders are delivered to you in the lab so you can determine their identity.

Due to limitations in equipment at the school, you have been asked to use a simple series of tests to determine the identity of the powders. Six known white powders have been provided. You will run tests on each of the six known powders and record your results. Next, you will test an unknown powder to determine its identity.

Materials: (you need to complete this list) White powder samples A,B,C,D,E,F Spot plate

Precautions:Sodium Carbonate Solution-

Acetic Acid-

Iodine Solution-

Procedure:Part A

1. Place a very small amount of sample A on the black paper and observe its appearance with a hand lens or microscope. Record appearance, draw a sketch if helpful. Repeat for other samples B-F.

2. Make a tray from aluminum foil with 6 sections to it. Place a small scoop of each sample on the tray. Draw a sketch in your notebook of the tray and the location of each white powder.

3. Place the foil tray on a hot plate set at a medium setting. Observe each substance for several minutes. Record your results carefully. Dispose of foil and contents in trash can.

4. Place three small scoops of sample A in a test tube. Add about 10 mL of water. Shake from side to side for several seconds. Does the substance dissolve? Float? Record observations. Repeat for samples B-F. Do not dispose of these samples. You will need them in the next step.

5. Using the six test tubes from step 4, add about 5 mL of sodium carbonate solution, NaCO3, to each test tube. Record observations. Wash out tubes. Use a test tube brush and detergent.

6. Place a small scoop of each sample in individual wells of a spot plate. Add 2 drops of iodine solution to each well. Record your results. Wash spot plate.

7. Again, place a small scoop of each sample in individual wells of a spot plate. Add 3 drops of acetic acid to each well. Record your results. Wash spot plate.

8. Show your instructor your observations for the 6 known substances. She will initial or stamp your notebook.

Part B1. Once Part A has been completed you will be issued an unknown from your instructor.

Record the unknown number and carry out all of the same tests as in part A.2. Record all your observations.3. Compare the results of the unknown to those of the known powders.4. Determine which white powders A-F are in your unknown. It may be just one or a

mixture of the powders. 5. Record your results in your instructor’s lab notebook. Your correct identification of

the unknown is worth 30/100 lab points.

powder appearance

heat water sodiumcarbonate

iodine aceticacid

A

B

62

C

D

E

F

unknown

Chromatography

Purpose: To separate mixtures using the technique of paper chromatography. To calculate retention factors and determine if the same dye is used in different mixtures.

Materials: Filter paper RulerPencilMarkers

Solvents: ethanol and waterPlastic wrapCandies or food dyestoothpicks

Precautions:Ethanol is flammable.

Procedure:1. Draw a line using pencil 1.00cm from the bottom of the filter paper. 2. Using sharpie markers, make dots on the line about 1.00cm apart. (or use a toothpick wet

with water to make dots of color from candy coatings) Use different colors. 3. Concentrate the color by going over each dot repeatedly.4. Pour ethanol into beaker to .50 cm. (use salt water solution if using candies)5. Place paper into beaker making sure dots do not get wet by solvent. The solvent should be

below the 1.00 cm line initially. It will travel up the paper and carry pigments with it. 6. Cover beaker with plastic wrap and watch until solvent level nears the top of the paper.

Remove paper from beaker and draw over the solvent line with a pencil. 7. Calculate the retention factor for each color spot for each marker (or candy) used.

Note that there are no units of Rf. They are naked numbers.

Rf= distance traveled by pigment in cm distance traveled by solvent in cm

8. Compare Rf values. Are any of the pigments the same? How do you know?

Measuring Plastics

Purpose: To determine the mass to volume ratio of 5 different plastics using proper techniques for measuring and calculating with significant figures.

Materials: lab balance calculator Centimeter ruler tank of water 5 plastics

Procedure: 64

1. Record the appearance of each plastic. Use a lab balance to measure the mass of each plastic. Record masses.Use a ruler to measure the length, width, and height of plastic #1 to 0.01 cm. Record all values.Repeat step #2 for plastics 2-5.Calculate the volume of each solid using length x width x height.Follow the rules for sig figs.Divide mass/volume for each plastic. Follow the rules for sig figs.

Data:

plastic Mass (g)

Length (cm)

Width (cm)

Height (cm)

Volume (cm3)

Density g/cm3

Clear

Cloudy

White

Gray

Black

Analysis:Record the accepted values for the density of each plastic as provided by your instructor.

Clear =__________Cloudy =__________White =__________Gray =__________Black =__________

Calculate your % error for each plastic.

Clear=

Cloudy=

White=

Gray=

Black=

Sources of Error:

Post Lab Questions:1. Write the definition for the word density.

2. Several units of length may be inches, centimeters, or miles. List 3 possible units for density.

3. Based on your lab calculations, explain how density can be used to help scientists identify one solid from another.

4. Do you think density can be used to identify liquids or gases? Explain.

5. The density of water is 1.00g/mL, which is equal to 1.00g/cm3. Compare each of the plastic’s densities to that of water to predict if the plastic will sink of float when placed in a tank of water. Write the color of each plastic in the correct column on the chart below.I think it will…

float sink

6. We will test your predictions as a class. How many of your predictions where correct? __________

7. Define the following terms:a. Intensive property-

b. Extensive property-

8. Explain why density is an intensive property of matter.

66

Percent Sugar in Bubble Gum Names: Problem: What percent of bubble gum is sugar?

Hypothesis:

Experiment:

Materials: 5 pieces of bubble gum containing sugar paper cup

balance

Procedure:1. Use a balance to determine the mass of a clean paper cup. Record the mass in your data table.2. Unwrap 5 pieces of bubble gum containing sugar and place them in the cup.3. Determine the mass of the cup and the gum. Record the mass in your data table.4. Each person in the group should chew a piece of gum to remove the sugar.5. After about 5 minutes, collect the chewed gum in the massed cup and wash your hands.6. Determine the mass of the cup and gum. Record it in your data table.7. Calculate the mass of sugar dissolved from the gum (original mass of gum – final mass of

gum). Record the answer in your data table.8. Calculate the percentage of sugar in the gum by dividing the mass of the dissolved sugar by

the mass of the un-chewed gum and multiply by 100. Record the answer in your data table.

Data:

Mass of Paper Cup

Mass of Cup + Gum

Mass of Unchewed

Gum

Mass of Cup +

Chewed Gum

Mass of Chewed

GumMass of sugar

Percent of Sugar

Conclusions:1. What is the percent of sugar? (Show calculations below.) __________________________

2. What is the molar mass of the sugar, C12H22O11? (Show your work below.) _______________________

3. Convert the mass of dissolved sugar to moles. (Show your work below.) _______________________

4. How many molecules of sugar are in the dissolved sugar? (Show you work below.) _______________

68

Chalk it up to Chemistry!

Goals• Visualize the concept of the mole• Gain experience in calculating grams and moles

The ActivityIn this activity, you will visualize the concept of the mole using a mole of chalk as a model. You will practice calculations of moles and grams, and end up with a better understanding of what a mole is and how chemists use it.

Materials for Each Group• Lab sheet and pencil• A chunk of chalk• Electronic balance (share- use the same one before and after)• Periodic Table/Ion Sheet• A calculator

SAFETYNo special safety considerations are required.

Instructions- Complete the following. Show all calculations. Remember units and sig-figs!

1. Take a chunk of chalk and measure its mass on the balance:

2. Go outside and draw something on the sidewalk.

3. Go back in class and weigh the unused chalk.

4. Based on the initial mass of the chalk and the mass at the end, calculate how many grams of chalk you left out on the sidewalk.

5. Write down the molecular formula of the chalk. Its chemical name is calcium carbonate.

6. Calculate the molar mass of the chalk.

70

7. How many moles of chalk did you leave on the sidewalk?

8. How many molecules of chalk did you leave on the sidewalk?

9. How many atoms of calcium did you use?

10. How many atoms of carbon did you use?

11. How many atoms of oxygen did you use?

Summary

A. How much does a mole of chalk weigh? __________________________________________________________

B. How many grams did you leave outside? __________________________________________________________

C. Is it (circle the correct answer): 1. less than a mole

2. equal to a mole3. more than a mole

D. Define a mole, and explain its importance: ______________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________________

72

The Molar Volume of Hydrogen Gas

From Royal Society of Chemistry Student Sheets—Classic Chemistry Experiments

Introduction/PurposeOne mole of any gas occupies the same volume when measured under the same conditions of temperature and pressure. In this experiment, the volume of one mole of hydrogen is calculated at room temperature and pressure.

Precautions Wear goggles. Hydrochloric Acid is corrosive. BE EXTREMELY CAREFUL WITH THE BURETTE! You break it you bought it!

Procedure 1. Clean a piece of magnesium ribbon about 3.5 cm long and weigh accurately. (This should weigh between 0.02

and 0.04 g; if not adjust the amount used.)

2. Measure 25 mL of dilute hydrochloric acid into the burette. Carefully add 25 mL of water on top of this.

3. Push the magnesium into the end of the burette so it will stay in position with its own tension.

4. Add 50 mL of water to a 250 mL beaker.

5. Quickly invert the burette into the water. If this is done quickly and carefully very little is lost. It is important that the liquid level in the burette starts on the graduated scale. If it is not on the scale; momentarily open the tap, this allows the level to drop). Clamp the burette vertically.

6. Take the burette reading (care: it is upside down!)

7. Observe the magnesium react as the acid diffuses downwards, wait until all the magnesium has reacted.

8. Note the new volume on the burette (care: it is upside down).

9. Record your results.

Observations/Data1. Record your observations for each step.

2. Record the mass of magnesium used and the volume of hydrogen produced.

______________________________________________________________________________The equation for the reaction is Mg(s) +2HCl(aq) MgCl2(aq) + H2(g)

__________________________________________________________________

Calculations- Show the set up for each of the following.

1. Convert the mass of Mg used to moles of Mg.

2. Convert the volume of H2 gas produced from milliliters to liters.

3. Divide the liters of hydrogen gas produced by the moles of magnesium used.

4. Compare your experimental value of liters of gas per mole to the accepted value of 22.4 L per mole by calculating your percent error.

74

Hydrate Lab

Purpose: To determine the percentage of water in magnesium sulfate heptahydrate within 1% of the theoretical value and to determine the percentage of water in an unknown hydrate.

Materials: hot plate lab balance Beaker insulating pad Watchglass beaker tongs scoop magnesium sulfate heptahydrate (MgSO4 *7H2O) unknown hydrate salt

Precautions: Magnesium Sulfate Heptahydrate-

Procedure:A. Percentage of Water in Magnesium Sulfate Heptahydrate1. Mass a clean dry beaker and watch-glass. Record the mass.2. Remove the watch-glass using beaker tongs. Zero the balance. Use a scoop to add 1-2 grams of magnesium sulfate heptahydrate to the beaker. Record the exact mass of the Magnesium sulfate heptahydrate.3. Use beaker tongs to transfer the beaker to a hot plate. Do not touch the beaker or Watch-glass with your hands. This may change the mass and increase your %error.4. Cover the beaker with the watch-glass. Heat the hydrate gently to avoid spattering the salt. You will observe moisture on the sides of the beaker and the bottom of the watch-glass. Continue heating until all the moisture is evaporated; the salt will change from a crystalline to a powder form.5. Use beaker tongs to carefully transfer the beaker/watch-glass to an insulating pad. Allow it to cool for at least 5 minutes. Don’t touch. 6. Use tongs again to transfer the beaker/watch-glass to a balance. Record the final mass.7. Empty the beaker contents into a paper towel and discard. 8. Calculate the percentage of water for this trial. Record it in your instructor’s lab notebook as well as your own. Your instructor will let you know if you need to do a second trial or if you can proceed to the unknown.

B. Percentage of Water in an Unknown Hydrate1. Obtain an unknown hydrate from your instructor. Record the letter or number of your unknown.2. Repeat the above procedure with the unknown and again calculate and report the % of water in the hydrate.

Data: A. Percentage of Water in Magnesium Sulfate Heptahydrate

Mass of beaker and watch-glass: _________________________/______________

Mass of magnesium sulfate heptahydrate:_______________________/______________

Mass of beaker, watch-glass, and magnesium sulfate heptahydrate:__________/________

Mass of beaker, watch-glass and magnesium sulfate anhydrate:___________/_______

Mass of water:______________________________

Show calculation for % of water:

Mass of water x 100= Mass of hydrate

B. Percentage of Water in Unknown Hydrate unknown #___________

Mass of beaker and watch-glass: _________________________/______________

Mass of hydrate:_______________________/______________

Mass of beaker, watch-glass, and unknown hydrate:__________/________

Mass of beaker, watch-glass and unknown anhydrate:___________/_______

Mass of water:______________________________

Show calculation for % of water:

Mass of water x 100= Mass of hydrate

Conclusion: Possible Sources of Error:

76

Penny Isotopes Lab

EACH PERSON TURNS IN HIS/HER OWN PAPER; write on back of this sheet except data table

points possible

student assessment

points earned

Heading (first and last name, date, period, lab group) 1Title 1Objective 2Data table (including unknown #) 2Calculations (work shown and explanation provided) 3Post-lab questions 3Conclusion (correct number of each type of penny with reason)

3

Total 15

Background:In 1982, the U.S. government changed the composition of the penny; the actual amount of copper per penny was decreased because the price of the raw copper had become more than 1¢. Thus, pre-1982 and post-1982 pennies have different masses.

In this activity, a mixture of pre- and post-1982 pennies will represent the atoms of a naturally occurring mixture of two isotopes of the imaginary element “coinium.” Using the pennies, you will simulate one way scientists determine the relative amounts of different isotopes present in a sample of an element. You will be given a sealed container that holds 10 pennies – some mixture of pre-1982 and post-1982 pennies. Your container will hold any particular atomic mixture of the two “isotopes” – 9 of one isotope and 1 of the other isotope, or 8 of one and 2 of the other, etc. Your task is to determine how many of each of the isotopes of the element coinium are present without opening the container.

Real copper has 2 naturally occurring isotopes:63Cu Mass = 62.9298 g Occurs in 69.09% of all Cu atoms65Cu Mass = 64.9278 g Occurs in 30.91% of all Cu atoms

Recall that the atomic mass (as recorded on the Periodic Table) is really a weighted average of the isotopes:63Cu: (62.9298 g)(69.09%) = 43.48 amu65Cu: (64.9278 g)(30.91%) = 20.07 amu

Therefore, the atomic mass of Cu = 43.48 + 20.07 = 63.55 amu

Materials and Equipment

Electronic balance 1 post-1982 penny1 pre-1982 penny 1 empty container and 1 sealed container with a mixture of 10 pre- and

post-1982 pennies

Procedure:Record the number of your sealed container (most likely your group number).Measure and record the masses of a pre-1982 penny, a post 1982 penny, the empty container, and the sealed container.Use math to determine the number of pre- and post-1982 pennies in your sealed container.

Post-lab Questions:What is an isotope?In what ways is the penny mixture a good analogy or model for actual element isotopes?In what ways is the analogy misleading or incorrect?Name at least one other familiar item that could serve as a model for isotopes. Explain.

Unknown # : _____Mass of pre-1982 pennyMass of post-1982 pennyMass of empty container (envelope)Mass of sealed container (envelope with pennies)

*include proper units and sig figs

78

Classification of Chemical Substances

Purpose: To classify 10 chemical substances as ionic, molecular, macromolecular, or metallic based on physical properties.

Materials: Unknown substances Hand lens Test tubes Hot plate

Scoop Spot plate Conductivity apparatus

Precautions: *Handle all chemicals carefully. Do not inhale, ingest, or touch. *Keep flammable materials away from flame. Do not point test tube at anyone while heating. *Use conductivity apparatus properly.

Procedure: Appearance Record appearance of each substance. Melting Point

1. Place a small sample of 3-4 substances on a foil tray and place on a cool hot plate.

2. Heat gently. Record observations. ( low mp-melts, high mp-no change, decomposes- turns brown or black, releases water)

3. Repeat #2 for all substances.4. Discard foil in trash.

Solubility and Conductivity 1. Place a small scoop of substance in a clean test tube. Add a dropper full of water to the tube. Shake from side to side to dissolve. Record observations. ( very soluble, slightly soluble, or not soluble)2. Transfer some of water solution to a spot plate. Test electrical conductivity.

Rinse tester with distilled water. It is not necessary to test conductivity for substances that do not dissolve. (strong conductor, weak conductor, not a conductor)

3. Repeat for each substance.

Classification- Using your chart of substances, select the best classification for each substance.

I= ionic MM=macromolecular

M= molecular Me= metal

Substance Appearance Heating Solubility in water

Conductivity Classification

Post-Lab Questions:

1. Discuss why some substances were easy to classify.

2. Which substances were more challenging to classify? Why?

80

Chemical Reactions LabPurpose: To observe chemical reactions and classify the type of reaction and write a balanced chemical equation for each reaction.

Materials: 2 test tubesspot platepipetteslab burner & matchessilver nitrate solnsodium chloride soln

lead (II) nitrate solnsodium carbonate solncobalt (II) chloride solncopper(II) sulfate solncopper wiresolid aluminum

Precautions: _______________________________________________________________________ _______________________________________________________________________

Procedure: A. Mix the following solutions together in a spot plate and observe and record any changes.

1. silver nitrate and sodium chloride :

2. sodium chloride and lead (II) nitrate:

3. sodium chloride and sodium carbonate:

4. cobalt (II) chloride and sodium chloride:

5. copper (II) sulfate and sodium carbonate:

6. copper (II) sulfate and cobalt (II) chloride:

B. Put the following substances together in a test tube:

7. copper wire and silver nitrate:

8. copper wire and sodium chloride :

9. solid aluminum and copper (II) sulfate :

C. Heat a sample of the following in a test tube in the flame of a lab burner:

10. copper (II) sulfate :

82

Structure of Compounds Part 1

ManganeseForm ox # color procedure_________________________________

Mn ______ _________1. Manganese metal

MnO4 ______ _________ 2. Add 5 drops potassium permanganate sol’n into a test tube.

MnO4-2 ______ __________3. Add 3 drops of 6 M NaOH and 1 drop

sodium sulfite.

MnO2 ______ __________4. In a clean test tube, add 5 drops of Potassium permanganate. Add sodium Sulfite until the color fades. After Several minutes observe the solid particles of MnO2.

Mn+2 ______ __________5. In a clean test tube, add 5 drops of Potassium permanganate and 1 drop of 6 M HCl. Add 5 drops of sodium sulfite. Observe Mn+2.

Follow the procedure step by step. Record the color of the substance at each step. Calculate the oxidation number of the selected element in each form.

Chromium

DO NOT POUR CHROMIUM SOLUTIONS IN THE SINK! PUT IN WASTE JAR UNDER FUME HOOD.

Form ox # color procedure_________________________________

Cr ______ __________6. Chromium metal.

K2Cr2O7 ______ __________7. Place a few crystals of potassium dichromate into a dry test tube.

CrO3 ______ __________8. Add 5 drops of concentrated sulfuric acid.

CrO4-2 ______ __________9. In a clean test tube, add 10 drops of

potassium dichromate sol’n. Add 3 drops of 6 M NaOH.

Cr+3 ______ __________10. In a clean test tube, add 10 drops of potassium dichromate sol’n and 1 drop of 6 M HCl. Add 5 drops of sodium sulfite. *******************************************************************

Nitrogen

Form ox # color procedure_________________________________

N2 ______ _________11. Air contains over 78% nitrogen. Breathe the air. Note the color, odor, and taste.

NO ______ __________12. Put 10 drops of 6M HNO3 into a test tube. Add

84

A small piece of copper.

HNO3 ______ __________13. In a clean test tube, add 5 drops of concentrated HNO3. Observe the color of the acid.

NO2 ______ __________14. Add a small piece of copper.

NH4Cl ______ __________15. In a clean test tube, add a pea-sized portion of solid ammonium chloride.

NH3 ______ __________16. Add a 5 drops of 6 M sodium hydroxide and cover the end of the tube with your thumb for 30 seconds; shake the tube from side to side (DO NOT LET THE SOLUTION TOUCH YOUR THUMB). Release your thumb and Carefully. Observe any odor.

Structure of Compounds Part 2Complete the following exercises in your lab notebook.

Create a chart in your lab notebook like the one shown below. Hold your book sideways so the columns can be wide.

Draw a dot diagram for each formula listed. Use the kits to create models of each.

Draw and color a diagram of what the model looks like. Use VSEPR Theory to predict the shape of the molecule.

Chemical Formula

Dot Diagram

model shape

polarity

1. H2 H:H linear

nonpolar

2. Cl2

^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^ 1. H2

2. Cl2

3. Br2

4. I2

5. HCl

6. HBr

7. HI

8. H2O

9. CH4

10. CCl4

11. Cl2O

12. NH3

13. O2

14. N2

15. HCN

16. CO2

17. C2H6

18. C2H4

19. C2H2

20. CHCl3

86

21. CH3OH

22. H2O2

^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^^Model sphere colors: Carbon= blackNitrogen= blueOxygen=redChlorine= green

Fluorine= also use greenBromine= orangeIodine= purple

1

Density of LiquidsPurpose : To determine experimentally the density of isopropanol, acetone, and water.

Materials:Erlenmeyer flaskBuretRingstandBuret clampIsopropanolAcetoneWaterBalancefunnel

Precautions:Both acetone and isopropanol are flammable. Keep away from open flames. Keep plastics away from acetone. Be careful with glass.

Procedure:1. Set up equipment as shown in diagram.2. Fill the buret with isopopanol using a funnel. 3. Mass the empty Erlenmeyer flask. 4. Dispense about 5.00 mL of liquid into the flask, record the actual volume.5. Remass flask and contents.6. Add liquid in 5.0 mL increments to the flask remassing and recording actual volume each

time until the volume of liquid in the flask is 25.00 mL.7. Repeat the above procedure with water and acetone.8. Calculate the density for each trial of each liquid.9. Graph the mass to volume ratio for each liquid on the graph provided.10. (volume on the horizontal axis, mass on the vertical)11. Color code each line and draw a key next to the graph.12. Submit your densities to your instructor (yellow book)13. Calculate your % error.14. List as many sources of error as possible.

AcetoneMass of beaker + liquid (g)

Mass of liquid(g) Volume (mL) Density (g/mL)

0.00 0.00 -------

2

IsopropanolMass of beaker + liquid (g)

Mass of liquid (g) Volume (mL) Density (g/mL)

0.00 0.00 -------

WaterMass of beaker + liquid (g)

Mass of liquid (g) Volume (mL) Density (g/mL)

0.00 0.00 -------

Graph mass vs. volume of acetone, ethanol, and water.

3

Calculate % error: Accepted values will be provided by your instructor. (show work!)

List Possible Sources of Error:

4

Physical & Chemical ChangesPurpose: To classify changes observed in lab as either physical or chemical changes.Materials: read the procedure and write a list of all materials you will need to complete

this lab.

Precautions: Use MSDS to find information for each chemical used.MagnesiumCupric sulfateIronSilver nitrateSodium chlorideHydrochloric acidCalcium chlorideAmmonium chloridePotassium permanganate (solution K)Procedure: Record all observations in your lab notebook.

1. Cut a piece of magnesium ribbon about 5 cm long. Note the color, luster, and flexibility of the metal. Holding one end with crucible tongs ignite the other with your lab burner. Caution! Do not look directly at the burning magnesium!! Compare the ash with the original metal.

2. Add 2 pipettes full of HCl into a test tube. Add a small scoop of zinc. Record your observations. When the reaction is proceeding vigorously place a burning splint down into the mouth of the test tube. Result? What gas do you think is produced?

3. Add 2 pipettes full of cupric sulfate solution into a test tube. Note the appearance of each substance before you put them together. Add a small scoop of iron filings. Wait 5 minutes then examine the tube contents carefully. Record observations.

4. Add a pipette full of silver nitrate to a test tube. Add 5 drops of sodium chloride solution. Observations? Take the tube and contents to station #5.

5. Set up a funnel with filter paper. Dampen the paper using water from a wash bottle. Place a beaker under the funnel. Pour the contents of the test tube into the funnel. Use the wash bottle to rinse the tube adding the additional water to the funnel. Unfold the filter paper and expose the residue in it to sunlight for several minutes. Observations? Discard the filtrate (liquid remaining in the beaker) into the sink and rinse.

6. Put a pipette full of water each into 2 separate test tubes. Add a scoop of ammonium chloride to the first tube and a scoop of calcium chloride to the second. Shake each tube from side-to-

5

side. Place your hand around the bottom of each tube. Add a second scoop of the substance to each tube. Record your observations.

7. View the writing on the index card while you look through the mystery oxide solution. Record your observations.

8. Squirt some solution K on a piece of paper towel. Watch it for about a minute. Record your observations.

change observations

1. burning Mg

2. Zn + HCl

3. CuSO4 + Fe

4a. AgNO3 + NaCl

5a. filter

b. sunlight

6. NH4Cl + H2O

6

CaCl2 + H2O

7. mystery oxide

8. sol’n K + paper

Post-Lab Classify each of the following as an example of a physical or chemical change. Give a supporting statement for each.

1. ________________________ Burning Mg ribbon.________________________________________________________________________________________________________________________________________________

2. ________________________ zinc and sulfuric acid________________________________________________________________________________________________________________________________________________

3. _________________________ cupric sulfate solution and iron________________________________________________________________________________________________________________________________________________

4. _________________________ HCl and silver nitrate solution________________________________________________________________________________________________________________________________________________

__________________________ centrifuge mixture ________________________________________________________________________________________________________________________________________________

5. __________________________ filtering mixture

7

_________________________________________________________________________

________________________________________________________________________ __________________________ sunlight on white precipitate

_______________________________________________________________________

6. ___________________________ ammonium chloride and water _______________________________________________________________________ _______________________________________________________________________ ___________________________ calcium chloride and water _____________________________________________________________________

_____________________________________________________________________

7. ___________________________ mystery oxide__________________________________________________________________

______________________________________________________________________

____

8. ___________________________ solution K on paper towel__________________________________________________________________

_______________________________________________________________________

_____

8

Density of Carbon Dioxide

We will investigate a reaction that makes carbon dioxide in your stomach when an antacid is taken for stomach distress. You will let the gas escape from the reaction tube and measure its volume after collecting the carbon dioxide by water displacement. You will also find the mass and calculate the density of carbon dioxide.Materials:Antacid tablet broken into 2 or 3 piecesTest tubeStopper, glass bend, and rubber tubing to fitTub of water for water displacement of gasFlask or bottle to collect the gas, 250 or 400 mLGraduated cylinder, 100 mLLaboratory balance sensitive to 0.01 g

Procedure:

1. Place 10 mL of water in the test tube.2. Mass the test tube of water and the tablet pieces together. Keep the tablet pieces dry. Support the test tube in a beaker.3. Fill the flask or bottle with water. Cover the mouth of the bottle and turn it upside down in the tub of water. Put the gas delivery tube up inside the top of the bottle or flask.4. Drop the tablet pieces into the test tube and immediately put the stopper that is attached to the glass bend and the rubber tubing.5. Collect the gas in the inverted flask or bottle by displacing the water.6. When no more CO2 is produced, hold the flask so the level of water is the same inside and out.7. Mark the level of water with a piece of masking tape.8. Invert the flask letting out the water and CO2. 9. Fill the flask with water to the level marked by the tape. Pour this water into a graduated cylinder. This is also the volume of CO2 produced by the reaction.10. Find the mass of the test tube and contents.

Calculations:1. Subtract the final mass from the original mass of the test tube and tablet pieces.

2. Why is this mass less than the original mass?

3. Calculate the density of the gas produced.

10

4. Determine your percentage of error using the accepted value for the density of CO2

provided by your instructor.

11

Flame TestsPurpose: To determine the color created by metallic ions in a flame and to use the technique of flame tests to determine the identity of an unknown cation.Materials:Wooden splints soaked in waterLab burnerSolids containing copper (Cu), sodium(Na), barium(Ba), potassium(K), strontium(Sr), lithium(Li), and calcium (Ca) ions. These will be compounds containing the metals.Spectra-glassesUnknown cation solids or solutions.

Precautions:Long hair must be tied back and flammable materials kept from flame.All chemicals should be handled with care, avoiding direct contact.Clean up spills or drips immediately!

Procedure:1. Get one sample at a time of the substances. 2. Dip a soaked wooden splint into one of the solids, then hold it in the outer

flame of a lab burner. 3. Record the color of the flame observed. You can dip the splint into the solid

again to see the color. If the splint begins to burn the color produced by the cation will be masked.

4. Repeat this procedure for each substance using a new splint (or the other end of a splint) for each test. DO NOT CONTAMINATE THE SAMPLES!

5. Record your results.6. Get an unknown from your instructor. Conduct a flame test and identify the

metal. Compare to known samples if necessary.7. Record your unknown results in your instructor’s lab notebook.

Data:cation color

Barium

Calcium

Copper

Strontium

Sodium

Lithium

Potassium

12

Unknown _____________________Color produced in flame______________________________Cation indicated_____________________________

Metals vs. NonmetalsPurpose: To determine physical and chemical properties of unknown elements from the Periodic Table and based on experimental results classify each as a metal, semi-metal, or nonmetal.

Materials:

Precautions:

Procedure:Physical Properties

1. Record the color of each element.2. Hammer a small sample of each element on the tray provided.

Indicate if the element is malleable, brittle, or somewhere in between. Clean up after yourself disposing of unusable samples in the trash can. 3. Use the electrical conductivity tester to determine whether each

substance conducts electricity or not. Record your observations.

Chemical Properties1. Using tongs, heat a sample of each element in the flame of a lab

burner for about 1 minute. Record observations.2. Place a small sample of each element in a separate well of a spot plate

and cover each with distilled water. Let stand for 5 minutes. Record observations.

Next, add 3 drops of phenolphthalein to each well. Record observations.

3. In separate wells of a spot plate place a small piece of each element. Cover with 3 M HCl. After several minutes record observations.

Observations:

Conclusion: Classify each element as a metal, semi-metal, or nonmetal. Give reasons for each.

Element Classification Reasons:

A

B

14

C

D

E

F

15

Experiment Template

Experiment Title: INSERT EXPERIMENT TITLE HERE




Hazards:Insert hazard warnings and/or list hazardous chemicals here for students to look

up for themselves.






16

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