vcl workbook (high school)

PRENTICE HALL

Needham, MassachusettsUpper Saddle River, New Jersey

Virtual ChemLabRecord Sheets

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Copyright © by Pearson Education, Inc., publishing as Pearson Prentice Hall, Upper Saddle River, New Jersey 07458. All rights reserved. Printed in the United States of America. This publication is protected by copyright, and permissionshould be obtained from the publisher prior to any prohibited reproduction, storage in a retrieval system, or transmissionin any form or by any means, electronic, mechanical, photocopying, recording, or likewise. Student worksheets may bereproduced for classroom use, the number not to exceed the number of students in each class. Notice of copyright mustappear on all copies. For information regarding permission(s), write to: Rights and Permissions Department.

ISBN 0-13-166227-9

1 2 3 4 5 6 7 8 9 10 07 06 05 04 03

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Contents

Overview . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . v

Installing Virtual ChemLab . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . vii

Getting Started . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . ix

Lab 1: Flame Tests for Metals . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 1

Lab 2: Names and Formulas of Ionic Compounds . . . . . . . . . . . . . . . . . . . . . . . . . . 4

Lab 3: Counting by Measuring Mass . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 6

Lab 4: Thomson Cathode Ray Tube Experiment . . . . . . . . . . . . . . . . . . . . . . . . . . . . 9

Lab 5: Millikan Oil Drop Experiment . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 11

Lab 6: Atomic Structure: Rutherford’s Experiment . . . . . . . . . . . . . . . . . . . . . . . . . 14

Lab 7: Atomic Emission Spectra . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 17

Lab 8: Photoelectric Effect . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 19

Lab 9: Diffraction Experiments . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 22

Lab 10: Electronic State Energy Levels . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 24

Lab 11: Pressure-Volume Relationship for Gases . . . . . . . . . . . . . . . . . . . . . . . . . . . . 27

Lab 12: Temperature-Volume Relationship for Gases . . . . . . . . . . . . . . . . . . . . . . . . 30

Lab 13: Derivation of the Ideal Gas Law . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 32

Lab 14: Ideal vs Real Gases . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 34

Lab 15: Investigation of Gas Pressure and Mass . . . . . . . . . . . . . . . . . . . . . . . . . . . . 37

Lab 16: The Specific Heat of a Metal . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 39

Lab 17: Heat of Fusion of Water . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 42

Lab 18: Heats of Reaction . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 44

Lab 19: Heat of Combustion . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 47

Lab 20: Enthalpy and Entropy . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 49

Lab 21: Electrolytes . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 51

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Lab 22: Precipitation Reactions: Formation of Solids . . . . . . . . . . . . . . . . . . . . . . . . 53

Lab 23: Identification of Cations in Solution . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 55

Lab 24: Qualitative Analysis . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 59

Lab 25: Study of Acid-Base Titrations . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 60

Lab 26: Acid-Base Titrations . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 63

Lab 27: Ionization Constants of Weak Acids . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 66

Lab 28: Analysis of Baking Soda . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 68

Lab 29: Molecular Weight Determination by Acid-Base Titration . . . . . . . . . . . . . . 70

Lab 30: Redox Titrations: Determination of Iron . . . . . . . . . . . . . . . . . . . . . . . . . . . . 72

Answers . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 75

© Pearson Education, Inc., publishing as Pearson Prentice Hall. All rights reserved.

iv Contents

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OverviewWelcome to Virtual ChemLab for Prentice Hall Chemistry, a set of realistic andsophisticated simulations covering topics in general chemistry. In theselaboratories, students are put into a virtual environment where they are free tomake the choices and decisions that they would confront in an actual laboratorysetting and, in turn, experience the resulting consequences. These laboratoriesinclude simulations of inorganic qualitative analysis, fundamental experimentsin quantum chemistry, gas properties, titration experiments, and calorimetry.This version of Virtual ChemLab is intended to be used in conjunction with thePrentice Hall Chemistry Virtual ChemLab Record Sheets, which contains 30laboratory assignments covering the topics of stoichiometry, atomic theory, gasproperties, thermodynamics, chemical properties, acid-base chemistry, andelectrochemistry. Many of these “virtual” laboratory assignments can also befound as “real” assignments in the Student Editior or accompanying LaboratoryManual, but there are also a significant number of new assignments, such asThomson’s cathode ray tube experiment and the Millikan oil drop experiment,which will provide students the opportunity to perform experiments notpreviously available to them.

The general features of the inorganic simulation include 26 cations that canbe added to test tubes in any combination, 11 reagents that can be added to thetest tubes in any sequence and any number of times, necessary laboratorymanipulations, a lab book for recording results and observations, and astockroom for creating test tubes with known mixtures, generating practiceunknowns, or retrieving instructor assigned unknowns. The simulation usesover 2,500 actual pictures to show the results of reactions and over 220 videosto show the different flame tests. In all, there are in excess of 1016 possibleoutcomes for these simulations.

The purpose of the quantum laboratory is to allow students to explore andbetter understand the foundational experiments that led to the developmentof quantum mechanics. Because of the very sophisticated nature of most ofthese experiments, the quantum laboratory is the most “virtual” of the VirtualChemLab laboratory simulations. In general, the laboratory consists of an opticstable where a source, sample, modifier, and detector combination can beplaced to perform different experiments. These devices are located in thestockroom and can be taken out of the stockroom and placed in variouslocations on the optics table. The emphasis here is to teach students to probea sample (e.g., a gas, metal foil, two-slit screen, etc.) with a source (e.g., a laser,electron gun, alpha-particle source, etc.) and detect the outcome with a specificdetector (e.g., a phosphor screen, spectrometer, etc.). Heat, electric fields, ormagnetic fields can also be applied to modify an aspect of the experiment. Asin all Virtual ChemLab laboratories, the focus is to allow students the ability toexplore and discover, in a safe and level-appropriate setting, the concepts thatare important in the various areas of chemistry.

The gas experiments included in the Virtual ChemLab simulated laboratoryallow students to explore and better understand the behavior of ideal gases,real gases, and van der Waals gases (a model real gas). The gases laboratorycontains four experiments each of which includes the four variables used todescribe a gas: pressure (P), temperature (T), volume (V), and the number of

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Overview v

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moles (n). The four experiments differ by allowing one of these variables to bethe dependent variable while the others are independent. The fourexperiments include (1) V as a function of P, T, and n using a balloon to reflectthe volume changes; (2) P as a function of V, T, and n using a motor drivenpiston; (3) T as a function of P, V, and n again using a motor driven piston; and(4) V as a function of P, T, and n but this time using a frictionless, masslesspiston to reflect volume changes and using weights to apply pressure. Thegases that can be used in these experiments include an ideal gas; a van derWaals gas with parameters that can be changed to represent any real gas; realgases including N2, CO2, CH4, H2O, NH3, and He; and eight ideal gases withdifferent molecular weights that can be added to the experiments to form gasmixtures.

The virtual titration laboratory allows students to perform precise,quantitative titrations involving acid-base and electrochemical reactions. Theavailable laboratory equipment consists of a 50 mL buret, 5, 10, and 25 mL pipets,graduated cylinders, beakers, a stir plate, a set of 8 acid-base indicators, a pHmeter/voltmeter, a conductivity meter, and an analytical balance for weighingout solids. Acid-base titrations can be performed on any combination of mono-,di-, and tri-protic acids and mono-, di-, and tri-basic bases. The pH of thesetitrations can be monitored using a pH meter, an indicator, and a conductivitymeter as a function of volume, and this data can be saved to an electronic lab bookfor later analysis. Asmaller set of potentiometric titrations can also be performed.Systematic and random errors in the mass and volume measurements have beenincluded in the simulation by introducing buoyancy errors in the massweighings, volumetric errors in the glassware, and characteristic systematic andrandom errors in the pH/voltmeter and conductivity meter output. These errorscan be ignored, which will produce results and errors typically found in highschool-level laboratory work, or the buoyancy and volumetric errors can bemeasured and included in the calculations to produce results better than 0.1% inaccuracy and reproducibility.

The calorimetry laboratory provides students with three differentcalorimeters that allow them to measure various thermodynamic processesincluding heats of combustion, heats of solution, heats of reaction, the heatcapacity, and the heat of fusion of ice. The calorimeters provided in thesimulations are a classic “coffee cup” calorimeter, a dewar flask (a betterversion of a coffee cup), and a bomb calorimeter. The calorimetric method usedin each calorimeter is based on measuring the temperature change associatedwith the different thermodynamic processes. Students can choose from a wideselection of organic materials to measure the heats of combustion; salts tomeasure the heats of solution; acids, bases, oxidants, and reductants for heatsof reaction; metals and alloys for heat capacity measurements; and ice for amelting process. Temperature versus time data can be graphed during themeasurements and saved to the electronic lab book for later analysis.Systematic and random errors in the mass and volume measurements havebeen included in the simulation by introducing buoyancy errors in the massweighings, volumetric errors in the glassware, and characteristic systematicand random errors in the thermometer measurements.


vi Overview

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Installing Virtual ChemLabLocate and run the program “Setup ChemLab” (which is located in theappropriate operating system folder—“Mac” or “Windows”) on the CD-ROMdrive then follow the prompts. The CD is not needed to run the program.

Important Installation Notes and Issues

1. The graphics used in the simulations require the monitor to be set to 24-bit true color (millions of colors). Lower color resolutions can beused, but the graphics will not be as sharp.

2. In the directory where Virtual ChemLab is installed, the user must alwayshave read/write/erase privileges to that directory and all directoriesunderneath. This condition is initially set by the installer, but this mayhave to be reset manually if the system crashes hard while runningVirtual ChemLab. This is generally only a problem with the OSXoperating system if multiple users login to the same OSX machine. Thiscan also be a problem with advanced Windows operating systems if theuser is not a Power User or higher.

3. When multiple users access the same installed software on a givencomputer, file ownership and read/write privileges becomes a seriousissue since Virtual ChemLab shares some files among users to a certaindegree depending on the installation. (a) In a direct access installation orwhen multiple users on a network drive share the database, all usersmust have complete read/write/erase privileges to the directory (andall directories underneath) where the database is stored. (b) In a webconnectivity installation, either (i) the same computer login must beused for all ChemLab users (so file ownership is the same for all databasefiles) or (ii) each user who creates a local ChemLab account (or newChemLab user) must use the same computer login as when the accountwas created in order to maintain file ownership consistency. This willonly be a problem with OSX machines and Windows operating systemsusing Restricted Users.

4. When using Virtual ChemLab under the Windows 2000 Professional orhigher operating systems, users must be, at a minimum, a Power User inorder for the program to have sufficient rights to run properly. Theprogram will run as a Restricted User but the fonts will be incorrectalong with other minor annoyances. In a server environment where aRestricted User is necessary, we suggest that a separate ChemLabAccount be setup, which gives the user Power User Status but onlygives the user access to the ChemLab software. This is a MacromediaDirector limitation.©

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5. Occasionally on all Windows platforms, the Virtual ChemLab installerwill fail to load and execute. This can be corrected by going towww.javasoft.com and downloading and then installing the most recent

version of the java runtime software. The Virtual ChemLab installationsoftware is a java based application.

6. QuickTime 5.0 or later is required for the software to run properly. Themost recent version of QuickTime can be obtained athttp://www.apple.com/quicktime/

7. When the simulation software has been installed on a Windows 2000Professional operating system, there is better performance and bettersystem stability when the Windows 2000 Support Pack 2 has beeninstalled.

8. For unknown reasons, on some machines the QuickTime videos willnot play properly if the system QuickTime settings are in their defaultstate. This can be corrected by changing the Video Settings inQuickTime to Normal Mode.

9. Printing in Virtual ChemLab does not work inside the OSX operatingsystem.

10. There are occasional spontaneous shutdowns of the software in OSX.There are no known causes for this, but it appears to be a MacromediaDirector issue.


viii Getting Started

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Getting StartedVirtual ChemLab is launched by clicking on the VCL icon located on the desktopor in the Start Menu for Windows operating systems or on the Dock for theMacintosh OS X operating system. After the startup, students will be broughtto a hallway containing three doors and a workbook sitting on a table. Clickingon the electronic workbook opens and zooms into the workbook pages wherestudents can select preset assignments that correspond to the assignments inthe Prentice Hall Chemistry Virtual ChemLab Record Sheets. The Previous and Nextbuttons are used to page through the set of assignments, and the differentassignments can also be accessed by clicking on the section titles located onthe left page of the workbook. Clicking on the Enter Laboratory button willallow the students to enter the general chemistry laboratory (see below), andthe Exit button is used to leave Virtual ChemLab.

From the hallway, students can also enter the general chemistry laboratoryby clicking on the General Chemistry door. Once in the laboratory, studentswill find five different laboratory benches that represent the five differentgeneral chemistry laboratories. Rolling the mouse over each of these laboratorybenches pops up the name of the selected laboratory. To access a specificlaboratory, click on the appropriate laboratory bench. While in the generalchemistry laboratory, the full functionality of the simulation is available, andstudents can explore and perform experiments as dictated by their instructorsor by their own curiosity. The Exit signs in the general chemistry laboratoryare used to return to the hallway.

Detailed instructions on how to use each of the five laboratory simulationscan be found in the user guides located in the Virtual ChemLab installationdirectory. These same user guides can also be accessed inside each laboratoryby clicking on the Pull-Down TV and clicking on the Help button.

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Lab 1: Flame Tests for MetalsPurposeTo observe and identify metallic ions using flame tests.

BackgroundThe characteristic yellow of a candle flame comes from the glow of burningcarbon fragments. The carbon fragments are produced by the incompletecombustion reaction of the wick and candle wax. When elements, such ascarbon, are heated to high temperatures, some of their electrons are excitedto higher energy levels. When these excited electrons fall back to lowerenergy levels, they release excess energy in packages of light calledphotons, or light quanta.

The color of the emitted light depends on its energy. Blue light is moreenergetic than red light, for example. When heated, each element emits acharacteristic pattern of light energies, which is useful for identifying theelement. The characteristic colors of light produced when substances areheated in the flame of a gas burner are the basis of flame tests for severalelements.

In this experiment, you will perform the flame tests used to identifyseveral metallic elements.

Procedure1. Start Virtual ChemLab and select Flame Tests for Metals from the list of

assignments. The lab will open in the Inorganic laboratory.

2. Enter the stockroom by clicking inside the Stockroom window. Onceinside the stockroom, drag a test tube from the box and place it on themetal test tube stand. You can then click on a bottle of metal ion solutionon the shelf to add it to the test tube. When you have added one metalion, click Done to send the test tube back to the lab. Continue doing thisuntil you have sent one test tube for each the following metal ions to thelab: Na�, K�, Ca2�, Ba2�, Sr2�, and Cu2�.

3. On the right end of the supply shelf is a button labeled Unknowns. Clickon the Unknowns button to create a test tube with an unknown. Nowclick on each of the following bottles on the shelf: Na�, K�, Ca2�, Ba2�,Sr2�, and Cu2�. Do not change the maximum and minimum on the leftside. Click Save. An unknown test tube titled Practice will show in theblue rack. Drag the practice unknown test tube from the blue rack toplace it in the metal stand and click Done. Now click on the Return to Labarrow.

Name ___________________________ Date ___________________ Class _____________

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Flame Tests for Metals 1

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4. When you return to the lab you should note that you have seven testtubes. You will use two of the buttons across the bottom, Flame andFlame w/ Cobalt (blue glass held in front of the flame.) A test tube must bemoved from the blue test tube rack to the metal test tube stand in orderto perform the flame test. You can drag a test tube from the blue rack tothe metal test tube stand to switch places with a test tube in the metaltest tube stand. Just above the periodic table there is a handle. Click onthe handle to pull down the TV monitor. With the monitor down youcan mouse-over each test tube and it will identify what metal ion the testtube contains. As you mouse over each test tube, you will also see apicture of what it contains in the lower left corner. One of your test tubesis labeled Practice and when you mouse over it, the TV monitor tells youit is an unknown.

5. Select the test tube containing Na� and place it on the metal stand. Clickthe Flame button. Record your observations in the data table below. Clickthe Flame w/Cobalt button and record your observations in the sametable.

6. Drag the K� test tube to the metal stand to exchange it with the Na�.Flame test K� with and without cobalt glass. Record your observationsin the table below.

7. For the other four ions, Flame test them only. Do not use cobalt glass.Record your observations in the table below.

Ion Flame Color

sodium, Na�

sodium, Na� (cobalt glass)

potassium, K�

potassium, K� (cobalt glass)

calcium, Ca2�

barium, Ba2�

strontium, Sr2�

copper, Cu2�

unknown #1

unknown #2

unknown #3

unknown #4

Name ___________________________ Date ___________________ Class _____________


2 Flame Tests for Metals

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Flame Tests for Metals 3

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8. Flame test the practice unknown. Determine which of the six metal ionsit most closely matches. You may repeat the flame test on any of the sixmetal ions if necessary. When you are confident that you have identifiedthe unknown, open the Lab Book by clicking on it. On the left page, clickthe Report button. On the right page, click on the metal ion that youthink is in the practice unknown. Click Submit and then OK. If all of theion buttons turn green you have successfully identified the unknown. Ifany turn red then you were incorrect. Flame test the practice unknownagain to correctly identify your metal ion. Click on the red disposalbucket to clear all of your samples.

Analysis and Conclusions1. The energy of colored light increases in the order red, yellow, green,

blue, violet. List the metallic elements used in the flame tests inincreasing order of the energy of the light emitted.

2. What is the purpose of using the cobalt glass in the identification ofsodium and potassium?


Lab 2: Names and Formulas of Ionic CompoundsPurposeTo observe the formation of compounds and write their names andformulas.

Procedure1. Start Virtual ChemLab and select Names and Formulas of Ionic Compounds

from the list of assignments. The lab will open in the Inorganiclaboratory.

2. Enter the stockroom by clicking inside the Stockroom window. Onceinside the stockroom, drag a test tube from the box and place it on themetal test tube stand. You can then click on the bottle of Ag+ ion solutionon the shelf to add it to the test tube. Click Done to send the test tubeback to the lab. Click on the Return to Lab arrow.

3. Place the test tube containing the Ag+ solution in the metal test tubestand. Click on the Divide button on the bottom (with the large redarrow) four times to make four additional test tubes containing Ag+.With one test tube in the metal stand and four others in the blue rack,click on the Na2S bottle located on the lab bench. You will be able toobserve what happens in the window at the bottom left. Record yourobservation in the table below and write a correct chemical formula andname for the product of the reaction. If the solution remains clear, recordNR, for no reaction. Drag this test tube to the red disposal bucket on theright.

4. Place a second tube from the blue rack on the metal stand. Add Na2SO4.Record your observations and discard the tube. Use the next tube butadd NaCl, and record your observations. Use the next tube but addNaOH, and record your observations. With the last tube add Na2CO3 andrecord your observations. When you are completely finished, click onthe red disposal bucket to clear the lab.

5. Return to the stockroom and repeat steps 2–4 for Pb2�, Ca2�, Fe3�, andCu2�. Complete the table below.

AnalyzeEach cell should include a description of what you observed when thereagents were mixed and a correct chemical formula and name for allsolutions which turned cloudy and NR for all solutions which remainedclear. Remember to include roman numerals where appropriate.

Name ___________________________ Date ___________________ Class _____________


4 Names and Formulas of Ionic Compounds

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Names and Formulas of Ionic Compounds 5

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Ag� Pb2� Ca2� Fe3� Cu2�

Na2S(S2�)

Na2SO4(SO4

2�)

NaCl(Cl�)

NaOH(OH�)

Na2CO3(CO3

2�)


Name ___________________________ Date ___________________ Class _____________


6 Mollusks, Arthropods, and Echinoderms

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Lab 3: Counting by Measuring MassPurposeDetermine the mass of several samples of chemical elements andcompounds and use the data to count atoms.

ProcedureStart Virtual ChemLab and select Counting by Measuring Mass from the list ofassignments. The lab will open in the Calorimetry laboratory.

Part 1, Measuring Metal1. Click on the Stockroom. Click on the Metals sample cabinet. Open the top

drawer by clicking on it. When you open the drawer, a petri dish willshow up on the counter. Place the sample of gold (Au) in the sampledish by double-clicking on it. Zoom Out. Double-click on the petri dish tomove it to the stockroom counter. Click the green arrow to Return to Lab.

2. Drag the petri dish to the spotlight near the balance. Click on the Balancearea to zoom in. Drag the gold sample to the balance pan and record themass in Table 1.

3. Click on the red disposal bucket to clear the lab after each sample. Repeatfor lead (Pb), uranium (U), sodium (Na) and a metal of your choosing.

Table 1

gold (Au) lead (Pb) uranium (U) sodium (Na) Your Choice

Mass(grams)

Molar Mass (g/mol)

Moles of each element

Atoms of each element

6 Counting by Measuring Mass


Name ___________________________ Date ___________________ Class _____________

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Counting by Measuring Mass 7

Analyze1. Calculate the moles of Au contained in the sample and enter into Table 1.

2. Calculate the atoms of Au contained in the sample and enter into Table 1.

3. Repeat steps 1 and 2 for the other metals and fill in the table. Clear thelaboratory when you are finished by clicking on the disposal bucket.

Part 2, Measuring Compounds1. Click on the Stockroom. Double-click on sodium chloride (NaCl) on the

Salts shelf. The right and left arrows allow you to see additional bottles.

2. Return to Lab. Move the sample bottle to the spotlight near the balancearea. Click on the Balance area to zoom in and open the bottle by clickingon the lid (Remove Lid). Drag a piece of weigh paper to the balance panand Tare the balance.

3. Pick up the Scoop and scoop out some sample; as you drag your cursorand the scoop down the face of the bottle it picks up more. Select thelargest sample possible and drag the scoop to the weigh paper until itsnaps in place which will place the sample on the paper. Record themass of the sample in Table 2.

4. Repeat steps 1-3 for table sugar (sucrose, C12H22O11), NH4Cl, C6H5OH(phenol), and a compound of your choice. Record the mass of eachsample in Table 2.

Table 2

NaCl C12H22O11 NH4Cl C6H5OH Your Choice

Mass(grams)

Molar Mass (g/mol)

Moles of each element

Atoms of each element


Analyze1. Calculate the moles of C12H22O11 contained in the sample and record

your results in Table 2.

2. Calculate the moles of each element in C12H22O11 and record yourresults in Table 2.

3. Calculate the atoms of each element in C12H22O11 and record yourresults in Table 2.

4. Repeat steps 1–3 for the other compounds and record your results inTable 2.

5. Which of the compounds contains the most total atoms?

Name ___________________________ Date ___________________ Class _____________


8 Counting by Measuring Mass

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Lab 4: Thomson Cathode Ray TubeExperimentPurposeTo duplicate the Thomson cathode ray tube experiment and calculate fromcollected data the charge to mass ratio (q/me) of an electron.

BackgroundAs scientists began to examine atoms, their first discovery was that theycould extract negatively charged particles from atoms. They called theseparticles electrons. In order to understand the nature of these particles, theywanted to know how much charge they carried and how much theyweighed. In 1897, John J. Thomson showed that if you could measure howmuch a beam of electrons were bent in an electric field and in a magneticfield, you could determine the charge to mass ratio (q/me) for the particles(electrons). Knowing the charge to mass ratio (q/me) and either the chargeon the electron or the mass of the electron would allow you to calculate theother. Thomson could not obtain either in his cathode ray tube experimentsand had to be satisfied with just the charge to mass ratio.

Procedure1. Start Virtual ChemLab and select Thomson Cathode Ray Tube Experiment

from the list of assignments. The lab will open in the Quantumlaboratory.

2. What source is used in this experiment? (The source is on the left. Drag your cursor over it to identify it.)

What type of charge do electrons have?

What detector is used in this experiment?

3. Turn on the Phosphor Screen. What do you observe?

4. Drag the lab window down and left and the phosphor screen windowup and right in order to be able to minimize overlap. Push the Gridbutton on the phosphor screen, and set the Magnetic Field to 30 �T.(Click the button above the tens place three times.) What happens to thespot from the electron gun on the phosphor screen?

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5. Set the Magnetic Field back to zero and set the Electric Field to 10 V. Whathappens to the spot from the electron gun on the phosphor screen?

Where should the signal on the phosphor screen be if the electric andmagnetic forces are balanced?

6. Increase the voltage of the Electric Field to move the spot severalcentimeters from the center. To make your measurements more accurate,move the spot until it aligns with a grid marking. What is the voltage?

What is the distance from the center that the spot has moved (in cm)?

7. Increase the magnetic field strength until the spot reaches the center ofthe screen. What magnetic field creates a magnetic force that balancesthe electric force?

Summarize your data.

8. In a simplified and reduced form, the charge to mass ratio (q/me) can becalculated as follows:

where V � the electric field in volts, d � the deflected distance fromcenter in cm, and B � magnetic field in �T.

What is your calculated value for the charge to mass ratio for an electron(q/me)?

The modern accepted value is 1.76 � 1011. Calculate your percent erroras follows:

% Error � 0 your value � accepted value 0

accepted value � 100

q>me � (5.0826 � 1012) # V # d>B2

deflected distance (d) electric field (V) magnetic field (B)

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Lab 5: Millikan Oil Drop ExperimentPurposeIf you want to know either the charge or the mass of an electron, you needto have a way of measuring one or the other independently. In 1909, RobertMillikan showed that he could make very small oil drops and depositelectrons on these drops (1 to 10 electrons per drop). He would thenmeasure the total charge on the oil drops by deflecting the drops with anelectric field. You will get a chance to repeat his experiments, and, using theresults from Lab 4, be able to experimentally calculate the mass of anelectron.

Procedure1. Start Virtual ChemLab and select Millikan Oil Drop Experiment from the

list of assignments. The lab will open in the Quantum laboratory.

2. What source is used in this experiment?

How does this source affect the oil droplets in the oil mist chamber?

3. The detector in this experiment is a video camera with a microscopiceyepiece attached to view the oil droplets. Click the On/Off switch(red/green light) to turn the video camera on. What do you observe onthe video camera screen?

What force causes the drops to fall?

The oil drops fall at their terminal velocity, which is the maximumvelocity possible due to frictional forces such as air resistance. Theterminal velocity is a function of the radius of the drop. By measuringthe terminal velocity (vt) of a droplet, the radius (r) can be calculated.Then the mass (m) of the drop can be calculated from its radius and thedensity of the oil. Knowing the mass of the oil droplet will allow you tocalculate the charge (q) on the droplet.

4. Measure the terminal velocity of a drop. Select a small drop that is fallingnear the center scale and Click the Slow Motion button on the videocamera window when the drop appears at the top of the window. Waituntil the drop is at a tick mark and start the timer. Let the drop fall for atleast two or more tick marks and stop the timer. Do not let the drop falloff the end of the viewing scope. Each tick mark is 0.125 mm. Record thedistance and the time in the data table below.

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5. Measure the voltage required to stop the fall of the drop. You now need tostop the fall of the drop by applying an electric field between the twovoltage plates. Click on the buttons on the top or bottom of the ElectricField until the voltage is adjusted and the drop stops falling. This shouldbe done while in slow motion. When the drop appears stopped, turnslow motion off and do some final adjustments until the drop has notmoved for at least one minute. Record the voltage, V, indicated on thevoltage controller.

6. Calculate the terminal velocity and record the value. Calculate the terminalvelocity, vt, in units of m�s using this equation:

, where d is the distance the drop fell in meters and t is the

elapsed time in seconds.

Each of the equations in steps 7–9 are shown with units and without.You will find it easier to use the equation without units for yourcalculations.

7. Calculate the radius (r) of the drop and record the value. With theterminal velocity, you can calculate the radius of the using this equation:

(9.0407 � 10�5 without units)

8. Calculate the mass of the drop and record the value. You can use theanswer from your previous answer for r to calculate the mass of the dropgiven the density of the oil. The final equation to calculate the mass is

(3439.0 r3 , without units)

9. The force due to gravity must be the same as the force due to theelectric field acting on the electrons stuck to the drop: qE � mg. Usingthis equation, calculate the total charge (Qtot) on the oil drop due to theelectrons using the equation:

� (9.81 � 10�2 m/V,without units)Qtot � Q(n) # e � (9.810 � 10�2C # kg�1 # J�1 # m>V

� (3439.0kg # m�3) # r3m � Voil

# roil � 4p>3 # r3 # 821kg # m�3

"vt,"vt �r � (9.0407 � 10�5m1>2 # s1>2)

vt �dt

Voltage Time Distance Drop (V, in volts) (t, in seconds) (d, in meters)

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where Q(n) is the number of electrons on the drop, e is the fundamentalelectric charge of an electron, m is the mass calculated in #9, and V isthe voltage.

This answer will provide the total charge on the drop (Qtot). Thefundamental electric charge of an electron (e) is 1.6 � 10�19 C(coulombs). Divide your total charge (Qtot) by e and round your answerto the nearest whole number. This is the number of electrons (Q(n)) thatadhered to your drop. Now divide your total charge (Qtot) by Q(n) andyou will obtain your experimental value for the charge on one electron.

10. Complete the experiment and calculations for at least three drops andsummarize your results in the data table.

11. Average your results for the charge on one electron. Calculate thepercent error by:

What is your average charge for an electron?

What is your percent error?

12. You will recall that in the Thomson experiment you were able tocalculate the charge-to-mass ratio (q/me) as 1.7 � 1011. Using this valuefor q/me and your average charge on an electron, calculate the mass ofan electron in kg.

What is your calculated value for the mass of an electron in kg?

% Error �0your answer � 1.6 � 10�19 0

1.6 � 10�19 � 100%

Total charge on drop

Terminal Velocity Radius Mass (Qtot, in Charge on one Drop # (v t, in m/s) (r, in meters) (m, in kg) Coulombs) electron (C )

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Lab 6: Atomic Structure: Rutherford’s ExperimentPurposeTo discover how the physical properties, such as size and shape, of anobject can be measured by indirect means and to duplicate the gold foilexperiment of Ernest Rutherford.

BackgroundA key experiment in understanding the nature of atomic structure wascompleted by Ernest Rutherford in 1911. He set up an experiment thatdirected a beam of alpha particles (helium nuclei) through a gold foil andthen onto a detector screen. According to the plum pudding atomic model,scientists thought electrons floated around inside a cloud of positivecharge. Based on this model, Rutherford expected that almost all the alphaparticles should pass through the gold foil and not be deflected. A few ofthe alpha particles would experience a slight deflection due to theattraction to the negative electrons (alpha particles have a charge of �2).Imagine his surprise when a few alpha particles deflected at all angles,even nearly straight backwards.

According to the plum pudding model there was nothing in the atommassive enough to deflect the alpha particles. About this he said “. . . almostas incredible as if you fired a 15-inch shell at a piece of tissue paper and itcame back and hit you.” He suggested the experimental data could only beexplained if the majority of the mass of an atom was concentrated in a small,positively charged central nucleus. This experiment provided the evidenceneeded to prove this nuclear model of the atom. In this experiment, you willmake observations similar to those of Professor Rutherford.

Procedure1. Start Virtual ChemLab and select Atomic Structure: Rutherford’s Experiment

from the list of assignments. The lab will open in the Quantumlaboratory.

2. The experiment will be set up on the lab table. Point the cursor arrow tothe gray box on the left side. What particles are emitted from this source?

What are alpha particles?

3. Point the cursor arrow at the base of the metal sample stand (in thecenter) and squeeze the mouse. What metal foil is used?

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4. Point the cursor arrow to the detector (on the right). What detector isused in this experiment?

5. Turn on the detector by clicking on the red/green light switch. Whatdoes the signal in the middle of the screen represent?

What other signals do you see on the phosphor detection screen?

What do these signals represent?

Click the Persist button (the dotted arrow) on the phosphor detectorscreen. According to the plum pudding model, what is causing thedeflection of the alpha particles?

Make an observation of the number of alpha particles hitting thephosphor detection screen.

6. Now, you will make observations at different angles of deflection. Clickon the gray lab table window to bring it to the front. Grab the phosphordetection screen by its base and move it to the spotlight in the top rightcorner. The Persist button should still be on. Describe the number of hitsin this spotlight position as compared to the first detector position.

7. Move the detector to the top center spotlight position at a 90-degreeangle to the foil stand. Describe the number of hits in this spotlightposition as compared to the first detector position.

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8. Move the detector to the top left spotlight position. Describe the numberof hits in this spotlight position as compared to the first detectorposition.

What causes the alpha particles to deflect backwards?

How do these results disprove the plum pudding model? Keep in mindthat there are 1,000,000 alpha particles passing through the gold foil atany given second.

Are the gold atoms composed mostly of matter or empty space?

How does the Gold Foil Experiment show that almost all of the mass ofan atom is concentrated in a small positively charged central atom?

Further InvestigationStudents often ask, “Why did Rutherford use gold foil?” The most commonresponse is that gold is soft and malleable and can be made into very thinsheets of foil. There is another reason, which you can discover for yourself.

1. Turn off the phosphor detection screen. Double-click the base of themetal foil sample holder. It will move the holder to the stockroomwindow. Click on the Stockroom to enter. Click on the metal sample boxon the top shelf. Click on Na to select sodium. Return to Lab.

2. Move the metal foil sample holder from the stockroom window back tothe center spotlight. Turn on the phosphor detection screen. Click Persist.Observe the number of hits with sodium compared to the number of hitswith a gold sample. Why would Rutherford choose gold foil instead ofsodium foil? Explain.

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Lab 7: Atomic Emission SpectraPurposeTo view atomic emission spectra and use a spectrometer to measure thewavelength. The wavelength will be used to calculate frequency andenergy.

ProcedureStart Virtual ChemLab and select Atomic Emission Spectra from the list ofassignments.The lab will open in the Quantum laboratory.

The Spectrometer is on the right of the lab table. The emission spectra is in the detector window in the upper right corner with a graph of theIntensity vs � (wavelength).

Analyze1. How many distinct lines do you see and what are their colors? Draw

what you see.

2. Click on the Visible/Full switch to magnify only the visible spectrum. Youwill see four peaks in the spectrum. If you drag your cursor over a peak,it will identify the wavelength (in nm) in the x-coordinate field in thebottom right corner of the window. Record the wavelength in the tablebelow for the four peaks in the hydrogen spectrum. (Round to wholenumbers.)

3. Each wavelength corresponds to another property of light called itsfrequency. Use the wavelength value of each of the lines to calculate itsfrequency given that where c � 2.998 1017 nm/s (2.998 108 m/s).The energy (E) of a quantum of light an atom emits is related to itsfrequency (�) by E � h�. Use the frequency value for each line and h � 6.63 � 10�34 J.s to calculate its corresponding energy.

Wavelength (nm) Frequency (1/s) Energy (J)

Line #1 (left)

Line #2

Line #3

Line #4 (right)

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4. Now, investigate the emission spectra for a different element, helium.Helium is the next element after hydrogen on the periodic table andhas two electrons. Do you think the emission spectra for an atom withtwo electrons instead of one will be much different than hydrogen?

5. To exchange gas samples, turn off the Spectrometer with the On/Offswitch in the top right corner. Double-click on the Electric Field to placeit on the stockroom shelf. Double-click on the Gas (H2) sample tube toplace it on the stockroom shelf.

6. Click in the Stockroom. Click on the Gases samples on the top shelf. Clickon the cylinder labeled He to select helium as the gas and it will fill thegas sample tube. If you point to the gas sample tube it should read He.

7. Return to lab. Drag the gas sample tube off the stockroom shelf. Whenyou select it, a white spotlight will appear indicating where you canplace the gas sample tube—place it there. Drag the Electric Field andplace it on the gas sample tube. Carefully click the button just above theleft zero on the volt meter and change the voltage to 300 V. Turn on theSpectrometer. Click the Visible/Full switch to convert to only the visiblespectrum.

8. Draw the visible spectrum for helium. Is it different from hydrogen?

9. Determine the wavelength (in nm), the frequency (in 1/s) and theenergy (in J) for the peak on the far right.

Wavelength (nm) Frequency (1/s) Energy (J)

Line (far right)

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Lab 8: Photoelectric EffectPurposeTo duplicate photoelectric effect experiments.

BackgroundThough Albert Einstein is most famous for E � mc2 and his work indescribing relativity in mechanics, his Nobel Prize was for understanding avery simple experiment. It was long understood that if you directed light ofa certain wavelength at a piece of metal, it would emit electrons. In classicaltheory, the energy of the light was thought to be based on its intensity andnot its frequency. However, the results of the photoelectric effectcontradicted classical theory. These inconsistencies led Einstein to suggestthat we need to think of light as being comprised of particles (photons) andnot just as waves.

Each wavelength corresponds to another property of light calledfrequency. You will use the wavelength (�) value in the experiment tocalculate the frequency (�) given that where c � 2.998 � 1017 nm/s(2.998 � 108 m/s). The energy (E) of a quantum of light an atom emits isrelated to its frequency (n) by E � h� where h (Planck’s constant) �6.63 � 10�34J-s.

Procedure1. Start Virtual ChemLab and select Photoelectric Effect from the list of

assignments. The experiment opens in the Quantum laboratory.

2. What source is used in this experiment and what does it do?

At what intensity is the laser set?

At what wavelength is the laser set?

Record the wavelength (in nm) in the data table. Calculate the frequency(in 1/s) and the energy (in J) using the equations given in theBackground section of this lab. Determine the color of the light byclicking on the Spectrum Chart (just behind the laser); the markersindicate what color is represented by the wavelength selected.

Which metal foil is used in this experiment?

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What detector is used in this experiment and what does it measure?

Turn on the detector by clicking on the red/green light switch. Whatdoes the signal on the phosphor screen indicate about the laser lightshining on the sodium foil?

3. Decrease the Intensity to 1 photon/second. How does the signal change?

Increase the Intensity to 1 kW. How does the signal change?

4. Change the Intensity back to 1 nW and increase the Wavelength to 600nm. What do you observe? Record the wavelength in the data table.

Determine the maximum wavelength at which emission of electronsoccurs in the metal.

What is the difference between intensity and wavelength?

Which matters in the formation of photoelectrons: intensity orwavelength?

wavelength (nm) frequency (1/s) energy (J) light color

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5. Click in the Stockroom. Click on the clipboard and select PhotoelectricEffect (2). Return to Lab. The intensity is set at 1 nW and the wavelengthat 400 nm. The detector used in this experiment is a bolometer. Turn onthe bolometer by clicking the red/green light switch. This instrumentmeasures the kinetic energy of electrons emitted from the metal foil. Youshould see a green peak on the bolometer detection screen. The intensityor height of the signal corresponds to the number of electrons beingemitted from the metal. The x-axis is the kinetic energy of thephotoemitted electrons. Zoom in on the peak by clicking and draggingfrom half-way up the y-axis to the right past the peak and to the x-axis(it will draw an orange rectangle).

6. Increase and decrease the Intensity, what do you observe?

Increase and decrease the Wavelength, what do you observe?

What is the maximum wavelength that ejects electrons from the sodiummetal?

From this experiment, explain why violet light causes photoemission ofelectrons but orange light does not.


Lab 9: Diffraction ExperimentsBackgroundIt has long been known that if you shine light through narrow slits that arespaced at small intervals, the light will form a diffraction pattern. Adiffraction pattern is a series of light and dark patterns caused by waveinterference. The wave interference can be either constructive (lightpatterns) or destructive (dark patterns). In this experiment, you will shine alaser through a device with two slits where the spacing can be adjusted andinvestigate the patterns that will be made at a distance from the slits.

Procedure1. Start Virtual ChemLab and select Diffraction Experiments from the list of

assignments. The lab will open in the Quantum laboratory.

2. What source is used in this experiment and for what reason?

What is the wavelength of the Laser?

What is the spacing of the two slits on the two slit device?

Draw a picture of the pattern displayed on the video screen.

3. Change the Intensity of the Laser from 1 nW to 1W.Does the intensity of the light affect the diffraction pattern?

Change the Slit Spacing to 1 �m. Observe the pattern displayed on thevideo screen as you change the slit spacing from 1 �m to 7 �m byone-micrometer increments. What can you state about the relationshipbetween slit spacing and diffraction pattern?

4. Increase the Wavelength of the Laser to 700 nm.What affect does an increase in the wavelength have on the diffractionpattern?

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22 Diffraction Experiments

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5. Decrease the Intensity on the Laser to 1000 photons/second. Click on thePersist button on the video camera to look at individual photons comingthrough the slits. Observe for one minute. What observation can youmake about this pattern as compared to the pattern from the continuousbeam of photons?

Decrease the Intensity to 100 photons/second. Observe for anotherminute after clicking Persist. At these lower intensities (1,000 and 100photons/second), there is never a time when two photons go throughboth slits at the same time. How can a single photon diffract?

From this experiment, what conclusions can you make about the natureof light?

6. Click in the Stockroom. Click on the clipboard and select Two-SlitDiffraction—Electrons. Return to Lab. What source is used in thisexperiment?

Draw a picture of the diffraction pattern shown on the Phosphor Screen.

How does this diffraction pattern compare to the diffraction pattern forlight?

7. Decrease the Intensity to 10 electrons/second. The pattern now buildsone electron at a time. Click on the Persist button and observe for oneminute.

Has the diffraction pattern changed? Why or why not?

How can an electron diffract if there is only one?

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Lab 10: Electronic State Energy LevelsPurposeTo understand the origins of Quantum Theory by using a spectrometer toobserve the emission spectrum of several gases.

BackgroundThe classical picture of atoms would allow electrons to be at any energylevel. According to this classical model, when electrons are excited andthen fall back down to the ground state, they emit light at all wavelengthsand the emission spectrum would be continuous.

In the 1800s scientists found that when a sample of gas is excited by analternating electric current, it emits light only at certain discretewavelengths. This allowed for the development of spectroscopy, which isused in the identification and analysis of elements and compounds. Eventhough scientists found spectroscopy very useful, they could not explainwhy the spectrum was not continuous. The explanation of this was left toNiels Bohr, a Danish physicist. Bohr proposed that energy levels ofelectrons are not continuous but quantized. The electrons only exist inspecific energy levels. Because of this quantization of energy, excitedelectrons can only fall to discrete energy levels.

This assignment illustrates the measurements that helped Bohr develophis quantum model, now known as Quantum Theory. It also illustratessome practical uses for this science. Mercury vapor is used in fluorescentlights and sodium vapor in street lighting.

You can separate the lines in the full region of an emission spectrum byusing an optical prism or a diffraction grating. A spectrometer is aninstrument designed to separate the emitted light into its componentwavelengths and plots the intensity of the light as a function ofwavelength.

Procedure1. Start Virtual ChemLab and select Electronic State Energy Levels from the list

of assignments. The lab will open in the Quantum laboratory.

2. The lab table will be set up with four items. What is the detector on theright?

What is the metal sample?

A heat source is used to heat the metal sample to high temperatures.What is the temperature of the heat source?

3. Turn on the Spectrometer by clicking on the red/green light switch. Clickon the Visible/Full switch to change the view to the visible spectrum.

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Click on the Lab Book to open it. If any students in a previous class havesaved spectra, highlight and delete them. Click on the Record button (reddot on the spectrometer window) to record this spectrum in the labbook. Click just after the spectra file name and type tungsten metal.What observations can you make about the emission spectrum forheated tungsten metal?

4. Turn the Spectrometer off with the On/Off switch. Click on the Stockroom.Click the clipboard on the right. Click on the preset lab #9 Photoemission– H2 and return to the lab by clicking on the Return to Lab arrow. Clickon the Visible/Full switch to change the view to the visible spectrum.Record this spectrum in the Lab Book. Click just after the spectra file nameand type hydrogen gas. What observations can you make about theemission spectrum for hydrogen gas?

5. To exchange gas samples, turn off the Spectrometer. Double-click on theElectric Field to place it on the stockroom shelf. Double-click on the Gas(H2) sample tube to place it on the stockroom shelf.

6. Click in the Stockroom. Click on the Gases sample on the top shelf. Clickon the cylinder labeled Ne to select neon as the gas and it will fill the gassample tube. If you point to the gas sample tube it should read Ne.

7. Return to Lab. Drag the gas sample tube off the stockroom shelf. Whenyou select it, a white spotlight will appear indicating where you canplace the gas sample tube—place it there. Drag the Electric Field andplace it on the gas sample tube. Carefully click the button just above thefar left zero on the volt meter and change the voltage to 300 V. Turn onthe Spectrometer. Click the Visible/Full switch to convert to only thevisible spectrum. Record this spectrum in the lab book and identify thislink with the name of the element typed after the blue link.

8. Continue with this same process until you have completed the followinggas samples: H2, He, Ne, Na, and Hg. You should have five spectrasaved in the lab book in addition to tungsten metal. Record yourobservations for each element. You can return to the lab book and clickon any of the spectra to view them again. Include in your observations acomparison for each element to the spectrum for heated tungsten metal.

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9. How do your observations of these gas emission spectra help confirmQuantum Theory?

Application1. Load the spectrum for heated tungsten metal. Tungsten metal is used in

incandescent light bulbs as the heated filament.

2. Load the spectrum for mercury from the lab book into the spectrometer.Examine the visible spectrum. Click the switch to change to full spectrum.What differences do you see when changing between visible and fullspectrum for mercury?

Mercury vapor is used in the fluorescent light tubes that you see at school andhome. The emitted light is not very bright for just the mercury vapor, butwhen scientists examined the full spectrum for mercury they saw what youobserved and recorded. There is an enormous emission in the ultravioletrange (UV). This light is sometimes called black light. You may have seen itwith glow-in-the-dark displays.

Scientists coat the inside of the glass tube of fluorescent light tubes with acompound that will absorb UV and emit the energy as visible light with allthe colors of the rainbow. All colors together create white light which is whyfluorescent light tubes emit very white light.

3. Load the spectrum for sodium. What distinct feature do you see in thesodium spectrum?

Astronomers are excited about cities changing from normal street lights tosodium vapor street lights because astronomers can easily filter out the peakat 589 nm and minimize light pollution.

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Lab 11: Pressure-Volume Relationship for GasesPurposeInvestigate the relationship between the pressure and volume of a gas.

BackgroundRobert Boyle studied the properties of gases in the 17th century. He noticedthat gases behave similarly to springs; when compressed or expanded, theytend to ‘spring’ back to their original volume. You will make observationssimilar to those of Robert Boyle.

The purpose of this experiment is to learn about the relationshipbetween pressure and volume of an ideal gas. In order to do this, you willchange the pressure or volume while the other is kept constant, and thenyou will observe what happens.

Procedure1. Start Virtual ChemLab and select Pressure-Volume Relationship for Gases

from the list of assignments. The lab will open in the Gases laboratory.

2. Note that the cylinder contains an ideal gas with a molecular weight of4 g/mol. There is 0.3 mol of gas at a temperature of 298 K, a pressure of1 atm, and a volume of 7.336 L. To the left of the pressure meter is a leverthat will decrease and increase the pressure as it is moved up or down;the digit is changed depending on how far the lever is moved. Digitsmay also be clicked directly to reach the desired number. You may wantto practice adjusting the lever so that you can decrease and increase thepressure accurately. Make sure the moles, temperature, and pressure arereturned to their original values before proceeding.

3. Click on the Lab Book to open it. If data from a previous student appears,delete it. Once back in the lab, click the Save button to start recordingyour data. Increase the pressure from 1 atm to 10 atm one atmosphere ata time. Click Stop. A blue data link will appear in the lab book. To keeptrack of your data links, enter ‘Ideal Gas 1’ next to the link. Reset the lab.

4. Zoom Out by clicking the green arrow next to the Save button. ClickReturn Tank on the gas cylinder. Click switch to change from Real Gasesto Ideal Gases. Choose Ideal 8 (Ideal 8 MW � 222 g/mol). Click on theballoon chamber to Zoom In. Set parameters to 0.3 mol of gas at atemperature of 298 K and a pressure of 1 atm (the volume should be7.336 L). Repeat the experiment with this gas labeling the data link as‘Ideal Gas 8.’

Pressure-Volume Relationship for Gases 27


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5. Zoom Out by clicking on the green arrow next to the Save button. Click onthe Stockroom and then on the Clipboard and select Balloon Experiment N2.Set the parameters to 0.3 mol of gas at a temperature of 298 K, and apressure of 1 atm (the volume should be 7.334 L). You may have to click onthe Units button for some of the variables to change it to the correct unit.Repeat the experiment with this gas labeling the data link ‘Real Gas N2.’

Analyze1. Click the data link for Ideal Gas 1. Click the Select All button in the

Data Viewer window that will pop up. If you are using Windows useCTRL-C, or on a Macintosh CMD-C to copy the data. Paste the data intoa spreadsheet and create a graph with volume on the x-axis and pressureon the y-axis. Also create a graph for your data from Ideal Gas 8 andReal Gas N2.

2. Based on your data, what relationship exists between the pressure andvolume of a gas (assuming a constant temperature)?

3. Look up a statement of Boyle’s Law in your textbook. Do your resultsfurther prove this?)

4. Complete the tables from the data saved in your lab book. Use only asampling of the data for pressure at 1, 3, 6, and 9 atm.

Ideal Gas 1 MW � 4 g/mol

Ideal Gas 8 MW � 222 g/mol

Volume (L) Pressure (atm) PV Product (P � V )


28 Pressure-Volume Relationship for Gases


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Pressure-Volume Relationship for Gases 29

Real Gas N2

5. What conclusions can you make about the PV product with Ideal Gas 1,MW � 4 g/mol?

How is the PV product affected using an ideal gas with a differentmolecular weight (Ideal Gas 8)?

6. How are your results affected using a Real Gas (N2)?

Going FurtherBased on the results of this lab, develop a hypothesis to explain whyweather balloons are only partially filled before they are released into theatmosphere. (These balloons can reach altitudes of 40,000 ft.)



Lab 12: Temperature-Volume Relationship for GasesPurposeInvestigate the relationship between the temperature and volume of a gas.

BackgroundCharles’ Law was discovered by Joseph Louis Gay-Lussac in 1802, basedon unpublished work done by Jacques Charles in about 1787. Charles hadfound that a number of gases expand to the same extent over the same 80degree temperature interval. You will be observing properties similar tothose that Charles studied.

The purpose of this experiment is to learn about the relationshipbetween temperature and volume of an ideal gas. You will do this bykeeping all variables constant except temperature and volume to observewhat happens.

According to the kinetic theory of gases, the average kinetic energy ofthe gas increases with the temperature. When gas is in a rigid container, thepressure (exerted by the gas on the walls of the container) increases as thetemperature increases.

Procedure1. Start Virtual ChemLab and select Temperature-Volume Relationship for Gases


2. Note that you are using an Ideal gas with a molecular weight of 4g/mol. There is 0.05 mol of gas at a temperature of 100°C, a pressure of 1atm, and a volume of 1.531 L. To the left of the temperature meter is alever that will decrease and increase the temperature as it is moved upor down; the digit is changed depending on how far the lever is moved.Digits may also be clicked directly to type in the desired number, or theycan be rounded by clicking on the R button. You may want to practiceadjusting the lever so that you can decrease and increase thetemperature accurately. Make sure the moles, temperature, and pressureare returned to their original values before proceeding.

3. Click on the Lab Book to open it. If data from a previous student appears,delete it. Once back in the lab, click the Save button to start recordingyour data. Increase the temperature from 100°C to 1000°C 100 degrees ata time. Click Stop. A blue data link will appear in the lab book. To helpkeep track of your data links, enter ‘Ideal Gas 1’ next to the link.

Analyze1. Click the data link for Ideal Gas 1. Click the Select All button in the Data

Viewer window. If you are using Windows, use CTRL-C or on a MacintoshCMD-C to copy the data. Paste your data into a spreadsheet and create agraph with volume on the x-axis and temperature on the y-axis.

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30 Temperature-Volume Relationship for Gases


2. Based on your data, what relationship exists between the pressure andvolume of a gas (assuming constant temperature)?

3. Look up a statement of Charles’ Law in your textbook. Do your resultsfurther prove this?

4. Using the spreadsheet, fit the data to a curve (linear fit). The lowestpossible temperature is reached when an Ideal Gas has zero volume.This temperature is the y-intercept for the plotted line. What is thistemperature?

5. Zoom Out by clicking on the green arrow next to the Save button. Clickon the Stockroom and then on the Clipboard and select Balloon ExperimentN2. Set the parameters to 0.05 mol of gas at a temperature of 100°C, anda pressure of 1 atm (the volume should be 1.531 L). You may have tochange the units for some of the variables. Repeat the experiment thatwas performed on the ideal gas by increasing the temperature from100°C to 900°C by increments of 100°C saving the data to the lab bookand label the link as ‘Real Gas N2.’ Plot the data and fit it to a linearcurve.

6. What lowest temperature did you find for the Real Gas (N2)?

Going FurtherInterpret the results of this lab in terms of the kinetic theory of matter as itapplies to gases.

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Deri

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aw

Lab 13: Derivation of the Ideal Gas LawPurposeDerive the Ideal Gas Law from experimental procedures and determine thevalue of the Universal Gas Constant (R).

BackgroundAn ideal gas is a hypothetical gas whose pressure, volume, andtemperature follow the relationship PV � nRT. No ideal gases actuallyexist, although all real gases behave similarly to ideal gases near roomtemperature and pressure. All gases can be described to some extent usingthe Ideal Gas Law, which is important in our understanding of how allgases behave. You will be deriving the Ideal Gas Law.

The state of any gas can be described using the four variables: pressure(P), volume (V), temperature (T), and the number of moles of gas (n). Eachexperiment in the Gases Simulation allows three of these variables (theindependent variables) to be manipulated or changed and then shows theeffect on the remaining variable (the dependent variable).

Procedure1. Start Virtual ChemLab and select Derivation of the Ideal Gas Law from the

list of assignments. The lab will open in the Gases laboratory.

2. Use the balloon experiment to describe the relationship betweenpressure (P) and volume (V). Increase and decrease the pressure usingthe lever to the left of the pressure to determine the effect on volume.What can you conclude about the effect of pressure on volume? Write amathematical relationship.

3. Use this same experiment to describe the relationship betweentemperature (T) and volume, by increasing and decreasing thetemperature. What can you conclude about the effect of temperature onvolume? Write a mathematical relationship.

4. Use this same experiment to describe the relationship between moles ofgas and volume, by increasing and decreasing the number of moles (n).What can you conclude about the effect of moles on volume? Write amathematical relationship.

32 Derivation of Ideal Gas Law


5. Since volume is proportional to inverse pressure, temperature andmoles we can combine these three proportions into one proportionshowing V proportional to 1/P, T, and n. Write one combinedproportion to show the relationship of volume to pressure, temperatureand moles.

6. This proportional relationship can be converted into a mathematicalequation by inserting a proportionality constant (R) into the numeratoron the right side. Write this mathematical equation and rearrange withP on the left side with V.

7. This equation is known as the Ideal Gas Law. Using data for volume,temperature, pressure and moles from one gas apparatus experiment,calculate the value for R with the units . (Show all work andround to three significant digits).

8. Using the relationship between atmospheres and mm Hg of 1 atm � 760 mm Hg, calculate the value for R with the units

. (Show all work and round to three significant digits).

9. Using the relationship between atmospheres and kPa of 1 atm � 101.3 kPa, calculate the value for R with the units .(Show all work and round to three significant digits).

L # atmK # mol

L # atmK # mol

L # atmK # mol

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Derivation of Ideal Gas Law 33


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Ideal

vs R

eal

Gases

Lab 14: Ideal vs Real GasesPurposeInvestigate how temperature and pressure changes affect ideal and realgases.

BackgroundIn the Ideal Gas Law lab previously completed, you learned how to derivethe Ideal Gas Law (PV � nRT) from observations about the relationshipsbetween volume, temperature, pressure and moles. You calculated thevalue for the Universal Gas Constant (R) for an ideal gas. In thisexperiment you will also calculate the Universal Gas Constant (R), but withboth ideal and real gases and at high and low temperatures and pressures.

Procedure1. Start Virtual ChemLab and select Ideal vs Real Gases from the list of

assignments. The lab will open in the Gases laboratory with Ideal Gas 1.

2. Click on the Units buttons to change the units to L or mL for volume,atm for pressure, and K for temperature.

To change the value of pressure, temperature or moles, the lever to theleft of each number can be used. The value will decrease and increase asthe lever is moved up or down; the digit is changed depending on howfar the lever is moved. Digits may also be clicked directly to reach thedesired number. Clicking to the left of the farthest left digit will add thenext place; for example, if you have 1.7 atm you can click left of the 1and enter 2 to make it 21.7 atm or click left of the 2 and enter 5 to make it521.7 atm.

The small R button in the upper left corner rounds the number. Clickingseveral times will round from ones to tens to hundreds. The green arrowto the left of the Save button will Zoom Out. Clicking Return Tank on thegas cylinder will return the tank to the rack and allow you to select adifferent gas. Clicking the gas chamber will Zoom In to allow you tochange parameters.

Be careful not to make the balloon so large that it bursts. If it does, clickthe red Reset button in the top right and then reset your units and valuesfor each parameter.

Remember that volume must be in L. If mL appears, you must convertto L in your calculations.

34 Ideal vs Real Gases


3. Complete the Data Table for the following gases and conditions (all with0.1 mol):a. Ideal gas at low T � 10 K, high T � 1000 K, low P � 1 atm, high P �

15 atmb. Methane gas (CH4) at low T � 160 K, high T � 400 K, low P � 1 atm,

high P � 15 atmc. Carbon dioxide gas (CO2) at low T � 250 K, high T � 1000 K, low P �

1 atm, high P � 15 atm

Analyze1. If PV � nRT then R � PV/nT. Complete the table for each experiment

above. Use four significant digits.

Gas V (L) P (atm) T (K) n (mol)

Ideal, low T, low P

Ideal, low T, high P

Ideal, high T, low P

Ideal, high T, high P

CH4, low T, low P

CH4, low T, high P

CH4, high T, low P

CH4, high T, high P

CO2, low T, low P

CO2, low T, high P

CO2, high T, low P

CO2, high T, high P

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Ideal vs Real Gases 35


2. Which gases and conditions show significant deviation from the idealvalue of R? Explain.

Gas Calculated R (1-atm/K-mol)

Ideal, low T, low P

Ideal, low T, high P

Ideal, high T, low P

Ideal, high T, high P

CH4, low T, low P

CH4, low T, high P

CH4, high T, low P

CH4, high T, high P

CO2, low T, low P

CO2, low T, high P

CO2, high T, low P

CO2, high T, high P

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36 Ideal vs Real Gases

Ideal

vs R

eal

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Lab 15: Investigation of Gas Pressure and MassPurposeInvestigate the relationship between the internal pressure of a gas and theapplied external pressure by placing weights on a frictionless, masslesspiston.

BackgroundAn understanding of pressure is an integral part of our understanding ofthe behavior of gases. Pressure is defined as the force per unit area exertedby a gas or other medium. The pressure of a gas is affected by manyvariables, such as temperature, external pressure, volume, number of molesof a gas, and other factors. This experiment will help you to become morefamiliar with pressure.

Procedure1. Start Virtual ChemLab and select Investigation of Gas Pressure and Mass


2. To investigate the relationship between mass added to the pistonto apply an external pressure, note that in this experiment Pint � Pmass � Pext, where Pint is the internal pressure or the pressureof the gas in the cylinder, Pmass is the pressure being exerted on thegas by adding weights to the piston, and Pext is the pressure beingexerted on the piston by the gas in the chamber.

3. Click the green Piston button to move the piston into position. Recordthe force (in tons) and the internal pressure (in psi) in the Data Table.

4. Click on the tenths place for force and add 0.5 tons of force to the piston.Record the force and internal pressure in the table. Repeat this for2.5 tons (the weight of a small car).

external pressure calculated internal internal pressure force (tons) (psi) pressure (psi) from meter (psi)

0

.5

2.5

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Investigation of Gas Pressure and Mass 37


Analyze1. You must now calculate the pressure being exerted by the 0.5 tons or,

Pmass. First, convert tons to psi (pounds per square inch).How many pounds is 0.5 tons?

The diameter of the piston is 15 cm. What is the radius (in cm)?

1 inch � 2.54 cm. What is the radius of the piston in inches?

The area of the circular piston is found by . What is the area ofthe piston in square inches (in2)?

The pressure exerted on the piston by the added mass in pounds persquare inch (psi) can be determined by dividing the mass in pounds bythe area in square inches. What is the pressure exerted by the addedmass in psi?

The internal pressure is the sum of the external pressure and the addedmass. What is the calculated internal pressure? Compare your calculatedanswer to the internal pressure meter answer. How do they compare?

2. Predict the internal pressure (in psi) when 2.5 tons are added.

How does your calculated answer compare to the internal pressuremeter when you add 2.5 tons of mass? Record your data in the table.

A � �r2

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38 Investigation of Gas Pressure and Mass


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Lab 16: The Specific Heat of a MetalPurposeDetermine the specific heat of a metal using a calorimeter.

BackgroundOn a sunny day, the water in a swimming pool may warm up a degree ortwo while the concrete around the pool may become too hot to walk onwith bare feet. This may seem strange since the water and concrete arebeing heated by the same source—the sun. This evidence suggests it takesmore heat to raise the temperature of some substances than others, which istrue. The amount of heat required to raise the temperature of 1 g of asubstance by 1 degree K is called the specific heat capacity, or specific heat, ofthat substance. Water, for instance, has a specific heat of 4.18 J/K g. Thisvalue is high in comparison with the specific heats for other materials, suchas concrete or metals. In this experiment, you will use a simple calorimeterand your knowledge of the specific heat of water to determine the specificheat of several metals.

Procedure1. Start Virtual ChemLab and select The Specific Heat of a Metal from the list

of assignments. The lab will open in the Calorimetry laboratory.

2. Click on the Lab Book to open it. Record the mass of Pb on the balance. Ifit is too small to read click on the Balance area to zoom in, record themass of Pb in the Data Table, and return to the laboratory.

3. Pick up the Pb sample from the balance pan and place the sample in theoven. Click the oven door to close. The oven is set to heat to 200°C.

4. The calorimeter has been filled with 100 mL water. The density of waterat 25°C is 0.998 g/mL. Use the density of the water to determine themass from the volume and record the volume and mass in the DataTable.

Make certain the stirrer is On (you should be able to see the shaftrotating). Click the thermometer window to bring it to the front andclick Save to begin recording data. Allow 20–30 seconds to obtain abaseline temperature of the water.

5. Click on the Oven to open it. Drag the hot lead sample from the ovenuntil it snaps into place above the calorimeter and drop it in. Click thethermometer and graph windows to bring them to the front again andobserve the change in temperature in the graph window until it reachesa constant value and then wait an additional 20–30 seconds. Click Stop inthe temperature window. You can click on the Accelerate button on theclock to accelerate the time in the laboratory. A blue data link will appearin the lab book. Click the blue data link and record in the Data Table thetemperature before adding the Pb and the highest temperature afteradding the Pb.

#

The Specific Heat of a Metal 39


6. Repeat the experiment with a metal sample of your choosing. Click thered disposal bucket to clear the lab. Click on the Stockroom to enter.Double-click the Dewar calorimeter to move it to the Stockroom counter.Click the metal sample cabinet. Click a drawer (the samples arealphabetically arranged), and select a sample by double-clicking andzoom out. Double-click on the petri dish with the selected sample tomove it to the Stockroom counter. Return to the laboratory.

7. Move the petri dish with metal sample to the spotlight next to thebalance. Click on the Balance area to zoom and make sure the balancehas been tared. Move the metal sample to the balance pan and record themass in the table. Return to the laboratory.

8. Double-click the calorimeter to move it into position in the laboratory.Click the oven to open the door. Move the metal sample from thebalance pan to the oven and click to close the oven door. Click above thetens place several times on the front of the oven to change thetemperature to 200°C. Fill the 100 mL graduated cylinder with water byholding it under the tap until it returns to the counter and then pour itinto the calorimeter. Turn on the thermometer. Click on the Graph andSave buttons. Move your metal sample from the oven to the calorimeter.Follow the procedures used with Pb to obtain the equilibriumtemperature. Record your observations in the table.

Analysis and Conclusions1. Determine the changes in temperature of the water ( Twater).

2. Calculate the heat (q) gained by the water using the following equation:, given Cwater� 4.184 J/(g°C)qwater � mwater

# ¢Twater# Cwater

¢

Pb your choice

mass of metal (g)

volume of water (mL)

mass of water (g)

initial temperature of water (°C)

initial temperature of metal (°C)

max temp of water � metal (°C)

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40 The Specific Heat of a Metal


3. Determine the changes in temperature of the Pb ( TPb).

4. Remembering that the heat gained by the water is equal to the heat lostby the metal, calculate the specific heat of lead.

and

5. Calculate the percent error in the specific heat value that youdetermined experimentally. Use the accepted value given by yourteacher.

6. Repeat the calculation of specific heat capacity for your choice of metalsand calculate the percent error. Use the accepted value given by yourteacher.

7. What assumptions have been made that might affect your answers?

8. The calorimeter would probably absorb some of the heat. If we considerthis source of error we can correct for it by modifying the specific heat ofwater to include the heat absorbed by the calorimeter. The corrected specific heat of water for the calorimeter is 5.224 . Calculate the most accurate specific heat for lead and the metal of your choosing.What is the percent error for each using your new calculation?

Jk # g

% Error �0your answer � accepted answer 0

accepted answer � 100

CPb �qmetal

(mmetal)(¢Tmetal)mPb

# ¢TPb# CPbqwater � �qmetal �

¢

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Heat

of

Fusio

nof

Wate

r

Lab 17: Heat of Fusion of WaterPurposeTo measure the heat of fusion of water using a calorimeter.

Procedure1. Start Virtual ChemLab and select Heat of Fusion of Water from the list of

assignments. The lab will open in the Calorimetry laboratory with abeaker of ice on the balance and a coffee cup calorimeter on the labbench.

2. Click on the Lab Book to open it. Record the mass of the ice on thebalance. If it is too small to read, click on the Balance area to zoom in andrecord the mass of ice in the Data Table.

3. 100 mL of water is already in the coffee cup. Use the density of water at25°C (0.998 g/mL) to determine the mass from the volume and record inthe Data Table. Make certain the stirrer is On (you should be able to seethe shaft rotating). Click the thermometer window to bring it to the frontand click Save to begin recording data. Allow 20–30 seconds to obtain abaseline temperature of the water.

4. Drag the beaker from the balance area until it snaps into place above thecoffee cup and pour the ice into the calorimeter. Click the thermometerand graph windows to bring them to the front again and observe thechange in temperature in the graph window until it reaches a constantvalue. Click Stop in the temperature window. Click on the Acceleratebutton on the clock to accelerate the time in the laboratory. A blue datalink will appear in the lab book. Click the blue data link and record inthe Data Table the temperature before adding the ice and the lowesttemperature after adding ice.

5. If you want to repeat the experiment, click on the red disposal bucket toclear the lab, click on the Stockroom, click on the clipboard and selectPreset Experiment 3, Heat of Fusion of Water.

volume of water in calorimeter (mL)

mass of water in calorimeter (g)

mass of ice (g)

initial temperature (°C)

final temperature (°C)

42 Heat of Fusion of Water


Analysis and Conclusion1. Calculate T for the water by T � |Tf – Ti|.

2. Calculate the heat (q) transferred from the water to the ice by q � mC Twhere the heat capacity (C) for water is 4.18 J/g°C and the mass is forthe water in the calorimeter. Convert to kJ/mol.

3. Convert the mass of ice to moles.

4. Calculate Hfus of ice (kJ/mol) by dividing the heat transferred from thewater by the moles of ice melted.

5. Compare your experimental value of Hfus of ice with the acceptedvalue of 6.01 kJ/mol and calculate the % error as:

6. What possible sources of error could there be in this laboratoryprocedure?

% Error �0your answer � accepted answer 0


¢

¢

¢

¢¢

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Heat of Fusion of Water 43


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Heats

of

Reacti

on

Lab 18: Heats of ReactionPurposeMeasure the heats of reaction for three related exothermic reactions and toverify Hess’s law.

BackgroundEnergy changes occur in all chemical reactions; energy is either absorbed orreleased. If energy is released in the form of heat, the reaction is calledexothermic. If energy is absorbed, the reaction is called endothermic. In thisexperiment, you will measure the amount of heat released in these threerelated exothermic reactions:

NaOH (s) Na� (aq) � OH� (aq) � H1

NaOH (s) � H� (aq) � Cl� (aq) H2O � Na� (aq) � Cl� (aq) � H2

Na� (aq) � OH� (aq) � H� (aq) � Cl� (aq) H2O � Na� (aq) � Cl�(aq) � H3

After determining the heats of reaction, you will analyze your data andverify the additive nature of heats of reaction.

ProcedureStart Virtual ChemLab and select Heats of Reaction from the list ofassignments. The lab will open in the Calorimetry laboratory.

Reaction 1

1. There will be a bottle of NaOH near the balance (move the plotwindow to see it). A weigh paper will be on the balance withapproximately 4 g NaOH on the paper. The calorimeter will be on thelab bench and filled with 200 mL water. Click the Lab Book to open it.Make certain the stirrer is On (you should be able to see the shaftrotating). In the thermometer window click Save to begin recordingdata. Allow 20–30 seconds to obtain a baseline temperature of thewater.

2. Drag the weigh paper with the sample to the calorimeter until it snapsinto place and pour the sample into the calorimeter. Observe thechange in temperature until it reaches a maximum and then record datafor an additional 20–30 seconds. Click Stop. Click on the Acceleratebutton on the clock to accelerate the time in the laboratory. A blue datalink will appear in the Lab Book. Click the data link and record theinitial and final water temperatures in the Data Table. If you need torepeat this part of the experiment, enter the Stockroom and selectPreset Experiment #6 on the clipboard.

¢S

¢S

¢S

44 Heats of Reaction


Reaction 2

1. Click the red disposal bucket to clear the lab. Click on the Stockroom toenter. Click on the clipboard and select Preset Experiment #5. Return tothe laboratory.

2. There will be a bottle of NaOH near the balance. A weigh paper will beon the balance with approximately 4 g NaOH on the paper. Thecalorimeter will be on the lab bench and filled with 100 mL water, andthere will be a beaker containing 100 mL of 1.000 M HCl on the labbench. In the thermometer window click Save to begin recording data.Allow 20–30 seconds to obtain a baseline temperature of the water.

3. Make sure the beaker of HCl is visible and drag it to the calorimeterand pour it into the calorimeter. The HCl and the water are at the sametemperature so there should be no temperature change. Now drag theweigh paper with the NaOH to the calorimeter until it snaps into placeand pour the sample into the calorimeter. It is important that the HClbe added first and the NaOH added second. Observe the change intemperature until it reaches a maximum and then record data for anadditional 20–30 seconds. Record the temperature before adding theHCl and after adding the NaOH.

Reaction 3


2. In the thermometer window click Save to begin recording data. Allow20–30 seconds to obtain a baseline temperature of the water. Pour thefirst beaker containing the HCl into the calorimeter and then pour thesecond beaker containing the NaOH into the calorimeter. Observe thechange in temperature until it reaches a maximum and then record datafor an additional 20–30 seconds. Record the initial and finaltemperatures.

Analysis and Conclusions

1. Determine the change in temperature, T, for each reaction. Recordyour results in the table.

¢

Parameter Reaction 1 Reaction 2 Reaction 3

Mass NaOH



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Heats of Reaction 45


2. Calculate the mass of the reaction mixture in each reaction. (To do this,first determine the total volume of the solution. Then calculate the massof the solution, based on the assumption that the added solid does notchange the volume and that the density of the solution is the same asthat of pure water, 1.0 g/mL.) Remember to add the mass of the solid.Record your results in the table.

3. Calculate the total heat released in each reaction, assuming that thespecific heat capacity of the solution is the same as that of pure water,

. Record the result in the table.

4. Calculate the number of moles of NaOH used in reactions 1 and 2where n = m/MW. Record the results in the table.

5. Calculate the number of moles of NaOH used in reaction 3 bymultiplying the volume of NaOH times the molarity (1.000 mol/L).Record the results in the table.

6. Calculate the energy released in kJ/mol of NaOH for each reaction andrecord the results in the table.

7. Show that the equations for reactions 1 and 3, which are given in theBackground section, add to give the equation for reaction 2. Include theenergy released per mole of NaOH in each equation.

# Mass of Total Heat Heat Released Rxn Rxn Mixture DT Released mol NaOH per mol NaOH

1

2

3

4.184JK # g

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46 Heats of Reaction


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Heat o

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Lab 19: Heat of CombustionPurposeMeasure the heat of combustion of sugar.

BackgroundThe heat of combustion is the heat of reaction for the complete burning ofone mole of a substance. Calorimetry experiments such as thedetermination of the heat of combustion ( Hcomb) can be performed atconstant volume using a device called a bomb calorimeter.

In this assignment you will calculate the heat of combustion of tablesugar (sucrose, C12H22O11). The calorimeter must also be calibrated by firstcombusting benzoic acid.

Procedure1. Start Virtual ChemLab and select Heat of Combustion from the list of

assignments. The lab will open in the Calorimetry laboratory with thebomb calorimeter out, disassembled, and a sample of benzoic acid in thecalorimeter cup on the balance. The balance has already been tared.

Calibration of the calorimeter

2. Click on the Lab Book to open it. Highlight and delete any data links leftby a previous student.

3. Record the mass of the benzoic acid sample from the balance. If youcannot read it, click on the Balance area to zoom in, record the mass, andreturn to the laboratory.

4. Double-click in this order to assemble the calorimeter: (1) the cup on thebalance pan, (2) the bomb head, (3) the screw cap, and (4) the bomb.Click the calorimeter lid to close. Combustion experiments can take aconsiderable length of time. Click the clock on the wall labeled Accelerateto accelerate the laboratory time.

5. Click the Bomb control panel and the Plot window to bring them to thefront. Click on the Save button to save data to the lab book. Allow thegraph to proceed for 20–30 seconds to establish a baseline temperature.

6. Click Ignite and observe the graph. When the temperature has leveledoff (up to 5 minutes of laboratory time), click Stop. A blue data link willappear in the lab book. Click the blue data link to view the collecteddata. Record in the Data Table the temperature before and after ignitionof the benzoic acid sample.

Combustion of sugar


¢

Heat of Combustion 47


8. Repeat steps 4–7 for sucrose. Record your observations in the table.

Calibration of calorimeter

1. Calculate T for the water in the benzoic acid combustion by T �|Tf – Ti|.

Calculate the moles of benzoic acid (MW � 122.13 g/mol). n � masssample/molecular weight

2. Hcomb for benzoic acid can be calculated by H � (Csystem T)/n ,where n is the moles of benzoic acid in the sample and Csystem is the heatcapacity of the calorimetric system. If the accepted value for the heat ofcombustion for benzoic acid is 3226 kJ/mol, calculate the heat capacity(Csystem) of the calorimetric system.

Heat of combustion of sucrose

1. Calculate T for the water by T � |Tf – Ti|.

2. Calculate the moles of sucrose in the sample (MWsucrose � 342.3 g/mol).

3. Hcomb for sucrose can be calculated by , wheren is the moles of benzoic acid in the sample and Csystem is the heatcapacity of the calorimetric system. Using the value you calculated forCsystem, calculate the heat of combustion for sucrose.

4. If the accepted value for the heat of combustion for sugar is 5639kJ/mol, calculate the percent error.

%Error �0your answer � accepted answer 0


¢Hcomb � (Csystem¢T)>n¢

¢¢

¢¢¢

¢¢

benzoic acid (C7H6O2) sucrose (C12H22O11)

mass of sample (g)



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Heat

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48 Heat of Combustion


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Enth

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Lab 20: Enthalpy and EntropyPurposeObserve and measure energy changes during the formation of a solutionand to describe and explain those changes in terms of entropy and enthalpy.

Procedure1. Start Virtual ChemLab and select Enthalpy and Entropy from the list of

assignments. The lab will open in the Calorimetry laboratory.

2. There will be a bottle of sodium chloride (NaCl) on the lab bench. Aweigh paper will be on the balance with approximately 2 g of NaCl onthe paper.

3. The calorimeter will be on the lab bench and filled with 100 mL water.Click the Lab Book to open it. Make certain the stirrer is On (you shouldbe able to see the shaft rotating). In the thermometer window click Saveto begin recording data. Allow 20–30 seconds to obtain a baselinetemperature of the water.

4. Drag the weigh paper with the sample to the calorimeter until it snaps inplace and pour the sample in the calorimeter. Observe the change intemperature until it reaches a maximum and then record data for anadditional 20–30 seconds. Click Stop. Click on the Accelerate button onthe clock to accelerate the time in the laboratory.) A blue data link willappear in the Lab book. Click the data link and record the initial andfinal water temperatures in the Data Table.

5. Click the red disposal bucket to clear the lab. Click on the Stockroom toenter. Click on the clipboard and select Preset Experiment #7 and repeatthe experiment with NaNO3. Record the initial and final temperatures inthe Data Table.

6. Click the red disposal bucket to clear the lab. Click on the Stockroom toenter. Click the clipboard and select Preset Experiment #8 and repeat theexperiment with NaCH3COO (NaAc). Record the initial and finaltemperatures in the Table.

Mixture T1 T2 ¢T

NaCl (s) � H2O (l)

NaNO3 (s) � H2O (l)

NaCH3COO � H2O (l)

Enthalpy and Entropy 49


AnalyzeUse your experimental data to answer the following questions in or belowyour data table.

1. Calculate T for each mixture. T � T2 – T1

2. An exothermic process gives off heat (warms up). An endothermicprocess absorbs heat (cools off). Which solutions are endothermic andwhich are exothermic? What is the sign of the change in enthalpy in eachcase?

3. Which solution(s) had little or no change in temperature?

4. When sodium chloride dissolves in water, the ions dissociate.

NaCl (s) Na� (aq) � Cl� (aq)

Write ionic equations, similar to the one above, that describe howNaNO3 and NaCH3COO each dissociate as they dissolve in water.Include heat as a reactant or product in each equation.

5. What is the sign of the Gibbs free energy ( G) for each process?

6. Consider the Gibbs-Hemholtz equation, G � H � T S. For eachdissolving process, substitute the sign of G and H into the equationand predict the sign for the entropy ( S). Does the sign for entropy seemlogical? Explain

7. If the sign for G is negative (spontaneous process) and the sign for Sis positive (more disorder) for both dissolving processes, how could onebe endothermic (positive H) and one be exothermic (negative H)? Isthere more to consider than just the dissolving process?

¢¢

¢¢

¢¢¢

¢¢¢

¢

S

¢¢

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Enth

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50 Enthalpy and Entropy


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Ele

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Lab 21: ElectrolytesPurposeClassify compounds as electrolytes by testing their conductivity in aqueoussolution.

Procedure1. Start Virtual ChemLab and select Electrolytes from the list of assignments.

The lab will open in the Titration laboratory.

2. Click inside the Stockroom. Double-click on three reagents to move themto the stockroom counter: NaCl, Na2CO3 (100%), and NaHCO3 (100%).Return To Lab.

3. For each salt, complete the following procedure: double-click the bottleto move it to the spotlight next to the balance. Click the Beakers drawerand place a beaker in the spotlight next to the salt bottle. Click on theBalance area to zoom in and open the bottle by clicking on the lid(Remove Lid). Pick up the Scoop and scoop up some salt by dragging thescoop to the bottle and then down the face of the bottle. Pick up thelargest sample possible and place it in the beaker. Zoom Out.

Move the beaker to the stir plate. Pick up the 25 mL graduated cylindernear the sink and hold it under the water tap until it fills. Pour the waterby dragging the cylinder to the beaker. Turn on the conductivity meterand place the conductivity meter probe into the beaker and record theconductivity in the Data Table. Double-click the salt bottle to place it backon the Stockroom counter. Place the beaker in the red disposal bucket.

4. When you have completed the three reagents, return to the Stockroom.Double-click on each bottle to return it to the shelf. Obtain three moresamples (two salts and one solution): KNO3, NH4Cl, NH3. Return to Lab.Use the procedure from #4 for NH4Cl, and KNO3.

For the solution complete the following procedure: Place a beaker on thespotlight left of the stir plate. Click on the Pipets drawer and double-clicka 25 mL pipet. Pick up the bottle of solution from the Stockroom shelf andpour into the beaker until the beaker is at least one-quarter full. Thesolution bottle will automatically go back to the Stockroom shelf. Clickthe pipet bulb (Fill Pipet) to fill the pipet. Place the used beaker in the reddisposal bucket. Drag a new beaker to the spotlight under the pipet andclick the pipet bulb to Empty Pipet. Drag the beaker to the stir plate.Place the conductivity meter probe into the beaker and record theconductivity in the Data Table.

5. Return to the Stockroom. Double-click each bottle to return them to theStockroom shelves. Obtain two more samples: HCl and HCN. Measurethe conductivity of each solution using the procedure above, and recordthe conductivity in the Table.

Electrolytes 51


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Ele

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Analyze1. Electrolytes are compounds that conduct electricity in aqueous solution.

Which compounds in your table are electrolytes? Which are notelectrolytes?

2. Would any of these electrolytes conduct electricity in the solid form?Explain.

3. Are these ionic or covalent compounds? Classify each compound in thegrid as ionic or covalent. For a compound to be an electrolyte, whatmust happen when it dissolves in water?

4. When an ionic solid dissolves in water, water molecules attract the ions,causing them to dissociate, or come apart. The resulting dissolved ionsare electrically charged particles that allow the solution to conductelectricity. The following chemical equations represent thisphenomenon.

NaCl (s) Na� (aq) � Cl� (aq)Na2CO3 (s) 2Na� (aq) � CO3

2� (aq)

Write a similar balanced chemical equation for each electrolyte in the datatable.

5. After examining the chemical reactions for the electrolytes, why doesNa2CO3 have a higher conductivity than all of the other electrolytes?

SS

NaCl Na2CO3 NaHCO3 KNO3 NH4Cl NH3 HCl HCN

conductivity

52 Electrolytes


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Pre

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Form

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f Solid

s

Lab 22: Precipitation Reactions: Formation of SolidsPurposeObserve, identify, and write balanced equations for precipitation reactions.

Procedure1. Start Virtual ChemLab and select Precipitation Reactions: Formation of Solids

from the list of assignments. Lab will open in the Inorganic laboratory.

2. React each of the cations (across the top) with each of the anions (downthe left) according to the table below using the following procedures:

a. Click in the Stockroom. Once inside the stockroom, drag a test tubefrom the box and place it on the metal test tube stand. Then click onthe bottle of Ag� metal ion solution on the shelf to add it to the testtube. Click Done to send the test tube back to the lab. Return to Lab.

b. Drag the test tube containing the Ag� from the blue rack to the metaltest tube stand. Click on the Divide button (the large red arrow) fourtimes to make four additional test tubes containing Ag�. With onetest tube in the metal stand and four others in the blue rack, click onthe Na2CO3 bottle on the reagent shelf, observe what happens in thewindow at the bottom left. Record your observation in the tableabove. If the solution remains clear, record NR, for no reaction. Dragthis test tube to the red disposal bucket on the right.

c. Drag a second tube from the blue rack to the metal stand. Add Na2S,record your observations and discard the tube. Continue with thethird, fourth and fifth tube, but add NaOH, Na2SO4, and NaClrespectively. Record your observations and discard the tubes. Whenyou are completely finished, click on the red disposal bucket to clearthe lab.

d. Return to the stockroom and repeat steps a through c for five testtubes of Pb2� and Ca2�.

AgNO3 (Ag�) Pb(NO)3 (Pb2�) Ca(NO3)2 (Ca2�)

Na2CO3 (CO32�) a f k

Na2S (S2�) b g l

NaOH (OH�) c h m

Na2SO4 (SO42�) d i n

NaCl (Cl�) e j o

Precipitation Reactions: Formation of Solids 53


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Analyze1. Translate the following word equations into balanced chemical

equations and explain how the equations represent what happens ingrid spaces a and g.

a. In grid space a, sodium carbonate reacts with silver nitrate to producesodium nitrate and solid silver carbonate.

b. In grid space g, sodium sulfide reacts with lead (II) nitrate to producesodium nitrate and solid lead (II) sulfide.

2. Write a word equation to represent what happens in grid space m.

3. What happens in grid space d? What other reactions gave similarresults? Is it necessary to write an equation when no reaction occurs?Explain.

4. Write balanced equations for all precipitation reactions you observed.

5. Write balanced net ionic equations for all precipitation reactions youobserved.

54 Precipitation Reactions: Formation of Solids


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Identific

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atio

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Solu

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Lab 23: Identification of Cations in SolutionPurposeIdentify the ions in an unknown solution through the application ofchemical tests.

BackgroundThe process of determining the composition of a sample of matter byconducting a chemical test is called qualitative analysis. Solutions ofunknown ions can be subjected to chemical tests and the results can becompared to the results given by known ions. By conducting theappropriate tests and applying logic, the identities of the ions present inan unknown solution can be determined.

In this experiment, you will observe several types of chemical reactionscommonly used as tests in qualitative analysis. These reactions include thecolor of a flame as the chemical is placed in the flame and the formation ofa precipitate (solid).

Procedure1. Start Virtual ChemLab and select Identification of Cations in Solution from

the list of assignments. The lab will open in the Inorganic laboratory.

2. Enter the stockroom by clicking inside the Stockroom window. Onceinside the stockroom, drag a test tube from the box and place it on themetal test tube stand. You can then click on a bottle of metal ion solutionon the shelf to add it to the test tube. When you have added one metalion, click Done to send the test tube back to the lab. Repeat this processwith a new metal ion. Continue doing this until you have sent one testtube for each of the following metal ions to the lab: Na�, K�, and aNa�/K� mixture. Fill one test tube with just water by clicking on thebottle of distilled water. Now click on the Return to Lab arrow.

3. When you return to the lab you should note that you have four test tubes.Just above the periodic table there is a handle. Click on the handle to pulldown the TV monitor. With the monitor down you can drag your cursorover each test tube to identify what metal ion the test tube contains, andyou will see a picture of what it looks like in the lower left corner.

Part 1, Flame Tests1. You will use two of the buttons across the bottom, Flame and Flame w/

Cobalt (blue glass held in front of the flame.) A test tube must be movedfrom the blue test tube rack to the metal test tube stand in order toperform the flame test. You can drag a test tube from the blue rack to themetal test tube stand to switch places with a test tube in the metal testtube stand.

Identification of Cations in Solution 55


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2. Flame test sodium ion only. Flame w/ Cobalt test sodium ion only. Recordyour observations.

3. Flame test potassium ion only. Flame w/ Cobalt test potassium ion only.Record your observations.

4. Flame and Flame w/ Cobalt test a mixture of sodium and potassium.Record your observations.

5. Flame test a blank (distilled water) with and without cobalt glass to get afeel for what it looks like with no chemicals other than water. Recordyour observations.

6. Return to the Stockroom. On the right end of the supply shelf is a buttonlabeled Unknowns. Click on the Unknowns label to create a test tube withan unknown. Now click on Na� and K�. On the left side make theminimum � 0 and maximum � 2. Click the Save button. An unknowntest tube titled Practice will show in the blue rack. Drag the practiceunknown test tube from the blue rack to place it in the metal stand andclick Done to send it to the lab. Return to Lab.

7. Flame test the Practice Unknown and determine if it contains sodium orpotassium or both or neither. Click on the Lab Book. On the left page,click the Report button, Submit, then OK. If the ion button is green, youcorrectly determined whether the ion was present or not. If the ionbutton is red you did not make the correct analysis. Click the reddisposal bucket to clear the lab. If you want to repeat with a newpractice unknown, return to the stockroom and retrieve it from the bluerack. When you are confident that you can make a correct determinationwith sodium and potassium, proceed to Part 2.

Part 2, Insoluble Chlorides1. Return to the Stockroom. In a new test tube, place three ions: Ag�, Hg2

2�,and Pb2�. (There is Hg2� and Hg2

2� on the shelf. Make sure you obtainHg2

2�.) Return to Lab. As you proceed with the chemical analysis watchthe TV screen to see the chemistry involved in the chemical reactions.

56 Identification of Cations in Solution


2. Move the test tube to the metal stand. Click the reagent bottle NaCl toadd chloride to the test tube. What observations can you make?

Click the Centrifuge button. What observations can you make?

Each of the three ions form insoluble precipitates (solids) with chloride.If the solution turns cloudy white it indicates that at least one of thethree ions is present. Now, we must determine which one.

3. Turn the heat on with the Heat button. Observe the TV screen. Whathappened? If you cannot tell, turn the heat on and off while observingthe TV screen.

With the heat turned on, click Decant. Drag your cursor over the new testtube in the rack. What appears on the TV screen? What appears in thepicture window?

This is the test for Pb2�. If heated, it is soluble. When cooled it becomesinsoluble.

4. Turn off the Heat. Click the NH3 bottle on the reagent shelf. What do youobserve?

Addition of ammonia produces a diammine silver complex ion which issoluble. The mercury produces a black solid. This is the test for mercury.

5. Centrifuge and then Decant to pour the silver ion into another test tube.Move the tube with the black mercury solid to the red disposal bucket.Move the tube containing silver back to the metal stand. Click the pH 4reagent bottle. What do you observe?

The silver ion is soluble as the diammine silver complex ion in pH 10and is insoluble as AgCl in pH 4. You can click alternately on each of thepH bottles to confirm this test for silver ion.

6. Return to the Stockroom and create an unknown with Ag�, Hg22�, and

Pb2�. Use minimum � 0 and maximum � 3. Return to the lab andcomplete the analysis. Report your results in the Lab Book and check todetermine if you can correctly identify these three ions. When you areconfident that you can correctly identify Ag�, Hg2

2�, and Pb2� proceedto Part 3.

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Identification of Cations in Solution 57

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Part 3, Selected Transition Metal Ions1. Return to the Stockroom. In a new test tube, place three ions: Co2�, Cr3�,

and Cu2�. Return to Lab. As you proceed with the chemical analysis watchthe TV screen to see the chemistry involved in the chemical reactions.

2. Move the test tube to the metal stand. Click the NaOH bottle on thereagent shelf. What observations can you make?

3. Click Centrifuge and Decant. What observations can you make as youdrag your cursor over each test tube?

This is the test for chromium. If the new test tube in the blue rack isgreen when decanted then chromium is present. You can confirm it byplacing the clear green test tube in the metal stand and clicking pH 10and then adding HNO3. What observations can you make?

4. With the test tube containing the precipitate in the metal stand, addNH3. What observations can you make?

5. Centrifuge and Decant. Add HNO3 to the tube in the metal standcontaining the precipitate. What observations can you make?

This is the confirmatory test for cobalt ion (Co2�).

6. Place the test tube from the blue rack which is the decant from step # 5in the metal stand. Add HNO3. What observations can you make?

This is the confirmatory test for copper.

7. Return to the Stockroom and create an unknown with Co2�, Cr3�, andCu2�. Use minimum � 0 and maximum � 3. Return to the lab andcomplete the analysis. Report your results in the Lab Book and check todetermine if you can correctly identify these three ions.

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58 Identification of Cations in Solution

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Qualita

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Analys

is

Lab 24: Qualitative AnalysisPurposeDevelop a systematic panel of chemical tests to identify an unknownsolution of eight metal cations.

BackgroundIn the Identification of Cations in Solution experiment, you learned how toidentify eight metal cations in three groups. In this experiment you willcomplete a qualitative analysis scheme using the information learned in theprevious experiment with an unknown solution containing one to eight ofthe ions.

Procedure1. Start Virtual ChemLab and select Qualitative Analysis from the list of

assignments. The lab will open in the Inorganic laboratory.

2. Click to enter the Stockroom. Create an unknown with Na�, K�, Ag�, Hg2

2�, Pb2�, Co2�, Cr3�, and Cu2�. Set minimum � 0 andmaximum � 8. This means that you could have only water or anynumber of the ions up to all eight. Click Save. Move the PracticeUnknown to the metal stand and Return to Lab.

3. Move the test tube to the metal stand and click the Divide button threetimes. Move the three new test tubes to the right end of the blue rack.These three tubes are duplicates of your unknown. If you make amistake and need to begin again, you can use one of these. Before youuse the last one make additional duplicates.

4. From your experience in the previous lab, you know how to analyze thethree different groups: flame tests, insoluble chlorides, and selectedtransition metals. Now that all three groups are combined you will stillanalyze the unknown as three separate groups. First, flame test for Na�

and/or K�. Remember, neither may be present and one of the other sixions will have a unique flame test that you have not seen before. Second,test for the insoluble chlorides. If you obtain a precipitate when addingNaCl then Centrifuge and Decant. The decant will contain the thirdgroup. Third, complete the analysis on the transition metals.

5. Report your results in the Lab Book and check to see if you correctlyidentified the presence or absence of each of the eight cations.

Qualitative Analysis 59

4573_PH_CHEM_Lab24_pp059 8/11/04 12:00 PM Page 59

Lab 25: Study of Acid-Base TitrationsPurposeObserve the changes that occur during the titration of a strong acid andstrong base.

BackgroundTitrations provide a method of quantitatively measuring the concentrationof an unknown solution. In an acid-base titration, this is done by deliveringa titrant of known concentration to an analyte of known volume. (Theconcentration of an unknown titrant can also be determined by titrationwith an analyte of known concentration and volume.) Titration curves(graphs of volume vs. pH) have characteristic shapes. By comparison, thegraph can be used to determine the strength or weakness of an acid orbase. The equivalence point of the titration, or the point where the analytehas been completely consumed by the titrant, is identified by the pointwhere the pH changes rapidly over a small volume of titrant delivery.There is a steep incline or decline at this point of the titration curve. In thisassignment, you will observe this titration curve by titrating the strong acidHCl with the strong base NaOH.

Procedure1. Start Virtual ChemLab and select Study of Acid-Base Titrations from the list

of assignments. The lab will open in the Titration laboratory.

2. Click the Lab Book to open it; if other students have left data linkshighlight and delete them.

3. The buret will be filled with NaOH. The horizontal position of the orangehandle is off for the stopcock. Click the Save button in the Buret Zoom Viewwindow. Open the stopcock by pulling down on the orange handle. Thevertical position delivers solution the fastest with three intermediate ratesin between. Turn the stopcock to one of the fastest positions. Observe thetitration curve. When the volume reaches 35 mL, double-click thestopcock to stop the titration. Click Stop in the Buret Zoom View. A bluedata link will be created in the lab book, click on it to view the data.

Analyze1. The beaker contains 0.3000 M HCl and the buret contains 0.3000 M

NaOH. Write a complete balanced equation for the neutralizationreaction between HCl and NaOH.

The following questions can be answered by examining the Plot andData Viewer windows.

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Stu

dy

of

Acid

-Base

Tit

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60 Study of Acid-Base Titrations


2. What were the pH and color of the solution at the beginning of thetitration?

3. What were the pH and color of the solution at the end of the titration?

4. Examine the graph of the pH vs volume. Sketch the shape of the titrationgraph of pH vs volume (blue line).

5. What happens to the pH around 25 mL?

6. What caused the change observed in Question 4?

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Study of Acid-Base Titrations 61


7. Examine the graph of the conductivity vs volume. Sketch the shape ofthe titration graph of conductivity vs volume (red line).

8. What happens to the conductivity during the titration?

9. What caused the change observed in Question 8?

Further InvestigationComplete a similar laboratory activity except using a polyprotic acid. Click in the Stockroom. Click on the clipboard and select Polyprotic AcidStrong Base.

1. What observations can you make about the graph of a titration with apolyprotic acid?

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62 Study of Acid-Base Titrations

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Lab 26: Acid-Base TitrationsPurposeStandardize a NaOH solution and measure the molarity of an unknownacetic acid solution by titration with standardized NaOH.

BackgroundTitrations provide a method of quantitatively measuring the concentrationof an unknown solution. In an acid-base titration, this is done by deliveringa titrant of known concentration to an analyte of known volume. (Theconcentration of an unknown titrant can also be determined by titrationwith an analyte of known concentration and volume.) Titration curves(graphs of volume vs. pH) have characteristic shapes. By comparison, thegraph can be used to determine the strength or weakness of an acid orbase. The equivalence point of the titration, or where the analyte has beencompletely consumed by the titrant, is identified as the point where the pHchanges rapidly over a small volume of titrant delivery. There is a steepincline or decline at this point of the titration curve. In this assignment, youwill determine the molarity of an unknown solution of NaOH by using aprimary standard, potassium hydrogen phthalate (KHP). You will then usea standardized solution of NaOH to determine the molarity of an unknownsolution of acetic acid.

ProcedurePart 1

1. Start Virtual ChemLab and select Acid-Base Titrations from the list ofassignments. The lab will open in the Titrations laboratory.

2. Click the Lab Book to open it. Click the Buret Zoom View window to bringit to the front. Click the Beakers drawer and place a beaker in thespotlight next to the balance. Click on the Balance area to zoom in, openthe bottle of KHP by clicking on the lid (Remove Lid). Drag the beaker tothe balance to place it on the balance pan; tare the balance. Pick up theScoop and scoop out some KHP; as you drag your cursor and the scoopdown the face of the bottle it picks up more. Select the largest samplepossible and drag the scoop to the beaker until it snaps in place whichwill place the KHP in the beaker (about 1 g). Unload the full scoop twiceinto the beaker so you have around 2.0 g and record the mass of thesample in the Data Table. Place the beaker in the spotlight outside thebalance and Zoom Out.

3. Drag the beaker to the sink and hold it under the tap to add a smallamount of water. Place it on the stir plate and add the calibrated pHmeter probe to the beaker. Add Phenolphthalein for the indicator.

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4. The buret will be filled with NaOH. Click the Save button in the BuretZoom View window so the titration data be saved. The horizontalposition of the orange handle is off for the stopcock. Open the stopcockby pulling down on the orange handle. The vertical position deliverssolution the fastest with three intermediate rates in between. Turn thestopcock to one of the fastest positions. Observe the titration curve.When the blue line begins to turn up, double-click the stopcock to turn itoff. Move the stopcock down one position to add volume drop by drop.

There are two methods for determining the volume at the equivalencepoint: (1) Stop the titration when a color change occurs. Click the Stopbutton in the Buret Zoom View. A blue data link will appear in the labbook. Click the blue data link to open the Data View window. Scrolldown to the last data entry and record the volume at the equivalencepoint in Data Table 1; OR (2) Add drops slowly through the equivalencepoint until the pH reaches approximately 12. Click the Stop button in theBuret Zoom View. A blue data link will appear in the lab book. Click theblue data link to open the Data View window. Click Select All button tocopy and paste the data to a spreadsheet. Plot the first derivative of pHvs. volume. The peak will indicate the volume of the equivalence pointsince this is where the pH is changing the most as volume changes.

Repeat at least two additional times, and record your data in the table.

The molecular weight of KHP is 204.22 g/mol.

Analyze1. Write a balanced chemical equation for the reaction of KHP and NaOH.

2. What is the average molarity of the unknown NaOH for your closestthree titrations?

Trial mass KHP (g) volume NaOH (mL) molarity NaOH (mol/L)

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64 Acid-Base Titrations


Part 2

1. Click in the Stockroom. Click on the clipboard and select Weak Acid StrongBase Unknown. The base is 0.3060 M NaOH and the buret has been filledwith it. The weak acid is an unknown concentration of acetic acid. 25 mLof acetic acid has been added to the beaker. The indicator isphenolphthalein and has already been added. The pH meter has beenturned on and calibrated.

2. Open your lab book and click the Save button on the Buret Zoom Viewwindow before starting the titration. Begin the titration by opening thestopcock on the buret. Observe what happens to the pH as the titrationproceeds. As the titration nears the end point, decrease the rate ofdelivery to the slowest rate. Immediately after the color of the indicatorchanges, stop the titration and read the amount of titrant that has beendelivered from the buret by clicking Stop in the Buret Zoom Viewwindow. Open the blue data link in the lab book and determine thevolume at the equivalence point using one of the methods described inPart 1. Record the volume in the table.

3. Move the used beaker to the disposal bucket and repeat the experimentat least two more times. Fill the buret with NaOH. Place a beaker in thespotlight left of the stir plate and fill the beaker half full with acetic acid(HAc). Open the Pipets drawer and double-click on a 25 mL pipet. Clickthe pipet bulb (Fill Pipet) to fill the pipet. Move the beaker to thespotlight right of the stir plate and drag a new beaker to the spotlightunder the pipet and click the pipet bulb to Empty Pipet. Drag the beakerto the stir plate. Place the pH meter probe in the beaker and addphenolphthalein indicator. Make certain the lab book is open and clickSave in the Buret Zoom View.

Analyze1. Write a balanced equation for the reaction between HAc and NaOH.

2. What is the average molarity of the unknown acetic acid solution?

volume NaOH molarity NaOH volume HAc molarity HAc Trial (mL) (mol/L) (mL) (mol/L)

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Lab 27: Ionization Constants of Weak AcidsPurposeMeasure ionization constants of weak acids such as bromocresol green (BCG).

Procedure1. Start Virtual ChemLab and select Ionization Constants of Weak Acids from

the list of assignments. The lab will open in the Titration laboratory.

2. Using the bottle of 0.1104 M NaOH, fill the buret. The bottle will snapinto position to pour and remain until buret is full. Click the buret toopen the buret window.

3. Click the Beakers drawer and place a beaker in the spotlight to the left ofthe stir plate. Fill the beaker half full with 0.1031 M HAc. Open the Pipetsdrawer and double-click on a 10 mL pipet. Click the pipet bulb (FillPipet) to fill. Move the beaker to the spotlight to the right of the stir plateand place a new beaker under the pipet and click the pipet bulb toEmpty Pipet. Move the beaker with 10 mL HAc and place it on the stirplate. Click the Stir Plate Switch to turn it on. Place the calibrated pHmeter probe in the beaker. Add Bromocresol Green to the beaker (double-click on the bottle labeled Bro G).

4. What are the color and pH of the solution?

5. The horizontal position of the orange handle is off for the stopcock.Open the stopcock by pulling down on the orange handle. The verticalposition delivers solution the fastest with three intermediate rates inbetween off. Turn the stopcock two stops down from the off position.When there is a color change, double-click the stopcock to turn it off. If itis necessary to repeat the experiment, do not forget to refill the buretwith NaOH and place the pH meter in the beaker on the stir plate.

6. What are the color and pH of the solution?

Continue to add NaOH as before. What is the final color of the solution?

66 Ionization Constants of Weak Acids


Analyze1. An acid-base indicator is usually a weak acid with a characteristic color.

Because bromocresol green is an acid, it is convenient to represent itscomplex formula as HBCG. HBCG ionizes in water according to thefollowing equation:

HBCG � H2O BCG� � H3O�

(yellow) (blue)

The Ka (the equilibrium constant for the acid) expression is

When [BCG�] � [HBCG], then Ka � [H3O�]. From the pH of thesolution the [H3O�] and Ka can be determined.

2. What color is an equal mixture of HBCG and BCG�? What is the pH atthe first appearance of this color?

3. What is the Ka for bromocresol green?

You’re the ChemistDesign and carry out an experiment to measure the Ka of bromocresolpurple and methyl orange.

Ka �3BCG�4 3H3O�4

3HBCG4

S

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Lab 28: Analysis of Baking SodaPurposeDetermine the mass of sodium hydrogen carbonate in a sample of bakingsoda using stoichiometry.

Procedure1. Start Virtual ChemLab and select Analysis of Baking Soda from the list of

assignments. The lab will open in the Titration laboratory.

2. The beaker has 1.5000 g of the impure solid NaHCO3 and is filled withwater to make a volume of 25.00 mL. The indicator methyl orange is inthe beaker. The calibrated pH meter is in the beaker, the graph windowis open. Click on the Lab Book to open it and delete any previous datasaved by another students. Click on the Buret Zoom View window andthe pH meter window to bring them to the front. Click the Save buttonon the graph window before starting the titration. Begin the titration byopening the stopcock on the buret. Observe what happens to the pH asthe titration proceeds. It is important that the volume increments andpH measurements near the equivalence point are small enough so theequivalence point can be determined as closely as possible. When thegraph starts to curve downward decrease the rate of delivery to theslowest rate. Immediately after the color of the indicator changes, stopthe titration by double-clicking the stopcock. Click the Stop button in theBuret Zoom View window. Click the blue data link in the lab book.Scroll down the Data Viewer window to the last volume entry in the leftcolumn and record the volume in the Table.

There are two methods for determining the volume at the equivalencepoint: (1) Stop the titration when a color change occurs. Click the Stopbutton in the Buret Zoom View. A blue data link will appear in the labbook. Click the blue data link to open the Data View window. Scrolldown to the last data entry and record the volume at the equivalencepoint in Data Table 1; OR (2) Add drops slowly through the equivalencepoint until the pH reaches approximately 2. Click the Stop button in theBuret Zoom View. A blue data link will appear in the lab book. Click theblue data link to open the Data View window. Click Select All button tocopy and paste the data to a spreadsheet. Plot the first derivative of pHvs. volume. The peak will indicate the volume of the equivalence pointsince this is where the pH is changing the most as volume changes.

mass unknown sample (g) volume HCl (mL) molarity HCl (mol/L)

68 Analysis of Baking Soda


Analyze1. Write a balanced chemical equation for the reaction between NaHCO3

and HCl.

2. Calculate the moles of HCl by multiplying the volume of HCl in litersand the molarity of HCl in mol/L. (Keep four significant digits in all ofthe calculations.)

3. The moles of HCl can be converted to moles of NaHCO3 using thecoefficients from the balanced equation. What is the mole to mole ratio ofHCl to NaHCO3? How many moles of NaHCO3 are present in the sample?

4. Calculate the grams of NaHCO3 by multiplying the moles of NaHCO3by the molecular weight of NaHCO3 (84 g/mol).

5. The percent NaHCO3 present in the sample can be calculated bydividing the mass of NaHCO3 from Question 4 by the mass of thesample from the Table and multiplying by 100. What is the percentNaHCO3?

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Lab 29: Molecular Weight Determinationby Acid-Base TitrationPurposeDetermine the molecular weight of a solid acid by titration methods.

BackgroundTitrations provide a method of measuring the concentration of anunknown solution. In an acid-base titration, a titrant of knownconcentration is added to an analyte of known volume. (The concentrationof an unknown titrant can also be determined by titration with an analyteof known concentration and volume.) Titration curves (graphs of volumevs. pH) have characteristic shapes. The graph can be used to determine thestrength or weakness of an acid or base. The equivalence point of thetitration, where the analyte has been completely consumed by the titrant, isidentified as the point where the pH changes rapidly over a small volumeof titrant delivery. There is a steep incline or decline in the titration curve atthis point of the titration curve. In this assignment, you will determine themolecular weight of an unknown acid powder by weighing the solid todetermine the mass and titrating to determine the moles of acid.

Procedure1. Start Virtual ChemLab and select Molecular Weight Determination by Acid-

Base Titration from the list of assignments. The lab will open in theTitration laboratory.

2. Click the Lab Book to open it. Click the Buret Zoom View window to bringit to the front. Click the Beakers drawer and move a beaker to thespotlight next to the balance. Click on the Balance area to zoom in, openthe bottle by clicking on the lid (Remove Lid). Drag the beaker to thebalance to place it on the balance pan; Tare the balance. Pick up the Scoopand scoop out some sample; as you drag your cursor and the scoopdown the face of the bottle it picks up more. Select the largest samplepossible and drag the scoop to the beaker until it snaps in place whichwill place the sample in the beaker (about 1 g). Record the mass of thesample in the Table. Place the beaker in the spotlight outside the balanceand Zoom Out.

3. Drag the beaker to the sink and hold it under the tap to add a smallamount of water. Place it on the stir plate and add the calibrated pHmeter probe to the beaker. Add Phenolphthalein for the indicator.

4. The buret will be filled with NaOH. The horizontal position of theorange handle is off for the stopcock. Click the Save button in the BuretZoom View window. Open the stopcock by pulling down on the orangehandle. The vertical position delivers solution the fastest with threeintermediate rates in between. Turn the stopcock to one of the fastest

70 Molecular Weight Determination by Acid-Base Titration


positions. Observe the titration curve. When the blue line begins to turnup, double-click the stopcock to turn it off. Move the stopcock down oneposition to add solution drop by drop.

There are two methods for determining the volume at the equivalencepoint: (1) Stop the titration when a color change occurs. Click the Stopbutton in the Buret Zoom View. A blue data link will appear in the labbook. Click the blue data link to open the Data View window. Scrolldown to the last data entry and record the volume at the equivalencepoint in Data Table 1. OR (2) Add drops slowly through the equivalencepoint until the pH reaches approximately 12. Click the Stop button in theBuret Zoom View. A blue data link will appear in the lab book. Click theblue data link to open the Data View window. Click Select All button tocopy and paste the data to a spreadsheet. Plot the first derivative of pHvs volume. The peak will indicate the volume of the equivalence pointsince this is where the pH is changing the most as volume changes.

Repeat at least two additional times, and record your data in the table.Do not forget to refill the buret with NaOH and place the pH meter inthe beaker and add indicator each time.

The concentration of the NaOH is 0.1961 M. The moles of acid arecalculated by multiplying the volume of the NaOH (in L) by themolarity of the NaOH. Dividing the mass of the acid sample by themoles will provide the molecular weight in g/mol.

5. What is the average molecular weight of your three closest answers?

6. Calculate your percent error by the following formula using only thevalues used in the average if you complete more than three trials:

% Error �highest answer � lowest answer

average � 100

volume NaOH molarity NaOH (g/mol)Trial mass of acid sample (mL) (mol/L) molecular weight

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Redox

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Lab 30: Redox Titrations: Determinationof IronPurposeDetermine the percent composition of iron in a sample using an oxidation-reduction titration with potassium permanganate.

BackgroundTitrations provide a method of quantitatively measuring the concentrationof an unknown solution. In an acid-base titration, this is done by deliveringa titrant of known concentration to an analyte of known volume. (Theconcentration of an unknown titrant can also be determined by titrationwith an analyte of known concentration and volume.) Reduction-oxidation(redox) titrations can also be used to measure concentrations. In redoxtitrations, voltages of the mixture of an oxidant and reductant can also bemeasured as the titration proceeds. All titration curves have characteristicshapes. The equivalence point of the titration, or the point where theanalyte has been completely consumed by the titrant, is identified by thepoint where the voltage changes rapidly over a small volume of titrantdelivered. There is a steep incline or decline at this point of the titrationcurve. In this assignment, you will observe this titration curve by titratingKMnO4 with FeCl2.

Procedure1. Start Virtual ChemLab and select Redox Titrations: Determination of Iron

from the list of assignments. The lab will open in the Titrationslaboratory.

2. Record the FeCl2 Unknown # in the Data Table. Click the Lab Book toopen it. Click the Buret Zoom View window to bring it to the front. Clickthe Beakers drawer and place a beaker in the spotlight next to thebalance. Move the FeCl2 bottle to the spotlight next to the balance. Clickon the Balance area to zoom in and open the bottle of FeCl2 by clickingon the lid (Remove Lid). Drag the beaker to the balance to place it on thebalance pan. Tare the balance. Pick up the Scoop and scoop up somesample by dragging the scoop to the bottle and then down the face ofthe bottle. Pick up the largest sample possible and place it in the beaker(about 1 g). Continue dragging the scoop to the beaker until it snaps inplace which will place the sample in the beaker. Unload the full scooptwice into the beaker so you have around 2.0 g and record the mass ofthe sample in the Table. Place the beaker in the spotlight outside thebalance and Zoom Out.

3. Place the beaker on the stir plate. Drag the 50 mL graduated cylinderunder the tap in the sink and fill it with distilled water. Pour the waterinto the beaker on the stir plate. Place the electrode in the beaker and turnon the volt meter.

72 Redox Titrations: Determination of Iron


4. The buret will be filled with KMnO4. The horizontal position of theorange handle is off for the stopcock. Click Save on the Buret window tosave the data to the lab book. Open the stopcock by pulling down on theorange handle. The vertical position delivers volume the fastest withthree intermediate rates in between. Turn the stopcock to one of thefastest positions. Observe the titration curve. When the blue line beginsto turn up, double-click the stopcock to turn it off. Move the stopcockdown one position to add volume drop by drop.

There are two methods for determining the volume at the equivalencepoint: (1) Stop the titration when a color change occurs. Click the Stopbutton in the Buret Zoom View. A blue data link will appear in the labbook. Click the blue data link to open the Data View window. Scrolldown to the last data entry and record the volume at the equivalencepoint in the Data Table; OR (2) Add drops slowly through theequivalence point until the voltage reaches approximately 1.50 V. Clickthe Stop button in the Buret Zoom View. A blue data link will appear inthe lab book. Click the blue data link to open the Data View window.Click the Select All button to copy and paste the data to a spreadsheet.Plot the first derivative of voltage vs. volume. The peak will indicate thevolume of the equivalence point since it is where the voltage is changingthe most as volume changes.

5. Sketch the graph of the titration below. Label the axes.

The reduction potential of Fe2� is 0.732 volts. The reduction potential ofMnO4 in acidic solution is 1.507 volts. If you titrate FeCl2 into KMnO4,what happens to the voltage of the solution as the titration starts andproceeds to the end?

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Repeat the titration at least two additional times recording data in theTable. Do not forget to refill the buret with KMnO4, place the voltagemeter in the beaker and add water each time.

The molecular weight of FeCl2 is 151.91 g/mol.

Unknown # _____

Analyze1. Write a balanced net ionic equation for the reaction in acidic solution of

FeCl2 and KMnO4 (Fe2� becomes Fe3� and MnO4 becomes Mn2�).

2. The moles of MnO4� can be determined by multiplying the volume of

MnO4� required to reach the endpoint multiplied by the molarity of the

MnO4�. What are the moles of MnO4

�?

3. The moles of FeCl2 can be calculated by using the mole ratio from thebalanced equation. What are the moles of FeCl2?

4. The mass of FeCl2 in the sample can be calculated by multiplying themoles of FeCl2 by the molecular weight of FeCl2. What is the mass ofFeCl2 in the sample?

5. The percent iron in the unknown sample can be determined by dividingthe mass of FeCl2 in the sample by the total mass of the unknownsample. What is the percent iron in your unknown sample?

6. What is the average percent iron in the unknown sample using your bestthree answers?

Trial mass FeCl2 (g) volume KMnO4 (mL) molarity KMnO4 (mol/L)

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74 Redox Titrations: Determination of Iron

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vcl workbook (high school)

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