CHEMISTRY 59-241

PHYSICAL CHEMISTRY

LABORATORY MANUAL

WINTER 2008

3rd Edition, v.1.0

DEPARTMENT OF CHEMISTRY & BIOCHEMISTRY

UNIVERSITY OF WINDSOR

TABLE OF CONTENTS

Emergency Procedures............................................................................................................ 2

Safety Regulations................................................................................................................... 4

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

Policy on Plagiarism................................................................................................................ 9

Student Contract...................................................................................................................... 12

Marking Scheme and Outline................................................................................................. 13

EXPERIMENTS

Experiment 1: Solid – Liquid Phase Diagram ............................................................... 15

Experiment 2: Determination of Ka Using the Conductance Method ........................... 21

Experiment 3: The Iodine Clock ....................................................................................... 29

Experiment 4: Thermodynamic Functions of a Galvanic Cell ....................................... 37

Experiment 5: Adsorption from solution ........................................................................... 43

EMERGENCY PROCEDURES

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Campus Police 4444

Fire Department 911 or pull wall alarm

Ambulance Dispatch 911

Medical Office 7002

Poison Control 9-800-268-9017

First Aid Kits Rooms 172-2, 175, and 274-2

Eye Baths are in each hallway, 173-3 (Lab E), 173-6 (Lab F)

Safety Showers are in each laboratory

EMERGENCY PROCEDURES

Discovery of a Fire

1. Shout, “Fire”. Turn off all equipment. Close the windows and doors as youleave the room.

2. Activate the nearest fire alarm.

3. Evacuate the building via the nearest exit. Do NOT use elevators.

4. Report the fire to Campus Police.

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1. Advise students to remain calm and to stay away from the fire.

2. Do not attempt to fight fires that cannot be easily handled.

3. If you put out a fire with a fire extinguisher, NEVER WALK AWAY. Back away and standby in case the fire ignites.

Sounding of Evacuation Alarm

1. It is MANDATORY for University buildings to be evacuated during any fire alarm.

2. Place all flammable materials into safety cabinets.

3. Shut off all heat sources including lit Bunsen burners.

4. Close all doors and windows.

5. Ensure any handicapped persons are given assistance.

6. Evacuate the building quickly by walking out the nearest exit. DO NOT use the elevators.(If smoke is encountered in a stairwell or corridor, use an alternate route.)

7. The Building Fire Plan Managers and Fire Wardens (Orange Vests) will assume lead rolesin building evacuation and direct you the Assembly Area.

8. DO NOT re-enter the building until the Fire Department or Campus Police authorize it.

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SAFETY REGULATIONSSafety is a matter of the greatest importance to everyone in the Department. The principaldanger is fire, but others (toxic fumes, sharp objects, etc.) are also of concern. GraduateStudents have a double responsibility of concern as laboratory workers and teaching assistants.Most safety procedures derive from common sense. The basic rule of safety is “If you areunsure of the consequences of an action, DON’T DO IT!” The following sets outregulations, which apply in teaching laboratories. Strict adherence to these is a MUST. Defaulters may be excluded from the lab and course.

One of the most common accidents to occur in the laboratory is a chemical spill. Any chemicalspilled on yourself or on your clothing must be washed off with LARGE AMOUNTS OF WATER.The incident MUST be reported to the TA. If acid is spilled on floors or laboratory benches, it mustbe neutralized immediately (solid sodium bicarbonate) and then cleaned up with water after thegeneration of all gas (carbon dioxide) ceases. Obtain the assistance of the TA.

The following rules are strictly enforced:

Cell phones or any other electronic device are not permitted during scheduledlaboratory time. If a student is observed using any electronic device, the devicewill be confiscated until the laboratory session has been concluded.“Turn it off orturn it over!”

Safety glasses must be worn at all times in the laboratory. Students refusing to wearsafety glasses (or shatterproof spectacles) will be refused permission to perform theexperiment. Never wear contact lenses in the laboratory.

10. It is mandatory that a laboratory coat be worn in the laboratory.

# Laboratory work is only permitted on assigned laboratory periods in the assignedlaboratory room and in the presence of a TA. Unauthorized experimental workconducted outside of stipulated laboratory hours is prohibited and forbidden bydepartmental regulations on grounds of safety. Severe disciplinary action will be takenagainst anyone attempting unauthorized experiments in the laboratory.

# Only those people directly involved in the laboratory experiment are allowed in thelaboratory. Visitors are not permitted to enter the laboratory.

# Smoking, eating, and drinking are not permitted in the laboratory.

# Extraneous items (coats, books, etc.) should remain in the area designated for coatsand schoolbags. These items are not permitted at the workbenches.

# No open-toed shoes or sandals may be worn. Students not wearing the proper footwearwill not be permitted to enter the laboratory and perform the scheduled experiment.

# Long hair is a fire hazard and must be tied back at all times.

# Perform all reactions with toxic or poisonous reagents in the fume hood.

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# Transfer harmful reagents in the fume hood to eliminate the dispersion of toxic fumesthroughout the lab. Handle all solid reagents with a spatula and wear gloves asprotection if necessary.

# Never deliberately purposefully inhale (smell) or taste any chemical. The propertechnique to identify an odour is to fan across the top of the container with your hand towaft very dilute vapours towards you.

# Never place your face directly over the top of a container which is being heated or pointit at a neighbour. The contents may “bump” and be violently ejected from the container. This can happen even after the heat has been removed.

# Never use a Bunsen burner for ANY purpose unless instructed to do so by the TA.

# Adhere to the rules when disposing glassware and reagents. Use the containerdesignated for glass when disposing broken glassware and sharp objects, such asPasteur pipets. Use ONLY the properly labelled waste containers for disposal of liquidand solid chemical waste. Never dispose of ANY chemical into the sink.

# Keep working space clean and free of apparatus and/or other materials. This isparticularly important when flammable materials are in use. Wipe up spilled materialsimmediately.

# If products need to be stored for the next lab period, place products and intermediates inthe labelled desiccators. Never store any chemicals in the student lockers!

# Notify the TA of any accident, cut, or burn no matter how trivial it appears.

Accidents and Injuries

In the event of a fire in the laboratory, turn off all gas, and shut down the fume hoods if possiblebefore leaving the room. Do not hesitate to shout “FIRE” or sound the building alarm for anysizable fire. Close the doors and get out of the building through the nearest exit.

If a student’s clothes catches on fire, lie down and roll over repeatedly to smother the flames. Donot run about, which includes running any distance for a safety shower. Use the safety shower orfire blanket if very close by.

Chemical spills on bench tops, fume hoods, etc., should be cleaned up with paper towels andwashed. Treatment with an appropriate neutralising reagent, if necessary, should be based onconsultation with the TA and/or lab co-ordinator.

Dangerous chemicals that come into contact with skin or clothing should be washed off, followedby prolonged washing for 5 to 10 minutes. Do not be concerned with neutralization, just wash. Ifa sizable quantity of corrosive material is on the clothes, prevent skin contact by quickly removingthe clothing.

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Two eye baths are available in the laboratory for washing out eyes. Be sure to ask for help, nomatter how slight the eye injury is.

Internal contact with chemicals can occur through a cut from broken glass, or by inhalation oringestion. Wash out cuts and allow free bleeding to occur for a few minutes. The laboratory co-ordinator, who is trained in First Aid, can treat cuts. If shattered glass was involved, inspect forfragments still in the wound before bandaging. Move into fresh air for inhalation problems and donot exert oneself until breathing has become normal. Report at once if any chemicals may haveentered the mouth, whether or not if any chemical has been swallowed. Delay will not make anynecessary treatment less unpleasant or less necessary.

Burns should be immersed in very cold water for about 5 minutes and, if necessary, dressingshould be applied afterwards.

In general, remember the locations of the alarm box, showers, hoses, and extinguishers in caseof emergencies.

Important: MSDS (Material Safety Data Sheets) for all chemicals used inthe labs can be found in a binder at the back of the lab by thebalances.

Conduct, House Rules, and Safety

Anyone who does not conduct themselves responsibly in the lab (i.e., who persistently exhibitsthoughtlessness or silly behaviour) may, at the discretion of the TA, be asked to leave thelaboratory, and the professor and laboratory coordinator will be notified. Students who are askedto leave will not be able to make up the experiment nor will they be able to submit a lab report forthat experiment. A grade of zero for that experiment will result.

Experiments should be completed by the stated closing hour. Experimental work conductedoutside of stipulated laboratory hours or in the absence of a TA is prohibited by departmentalregulations on grounds of safety. Students are not competent and experienced enough to judgethe hazards, or lack of them, that may permit them to overrule this decision.

Good laboratory etiquette (e.g. noxious fumes contained in the hood, slippage and drippagemopped up, etc.) is expected from anyone working in the lab. Reasonably tidy bench tops andlockers are a sign of ordered activity and also make matters easier in case of water floods, benchfires, etc. Burners not in use must be turned off as their invisible flames are a distinct hazard.

The University of Windsor has a Policy and Procedure on Sexual Harassment. The followingstatement is drawn from the policy:

“The University of Windsor is committed to providing an environment for study, teaching, researchwork, and play for all members of the University community that is supportive of professional andpersonal development and free from sexual harassment.”

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SAFETY QUIZ

Please take a few moments to answer these questions.

1. Where is the location of the fire alarm pull?

___________________________________________________________________________

2. How many fire extinguishers are in the lab? Where is its/their location(s)?

___________________________________________________________________________

3. Where is the nearest eyewash?

___________________________________________________________________________

4. Where is the safety shower in the lab?

___________________________________________________________________________

5. Where is the fire blanket in the lab?

___________________________________________________________________________

Circle the correct answer to each question.

6. What type of footwear is required in the lab?(a) Dress shoes (b) Sandals(c) Shoes that cover toes and protect feet from spills

7. The main routes of exposure to chemicals are:(a) Eyes, nose, mouth and ears (b) Inhalation, ingestion and skin/eye contact(c) Through open-toed shoes

8. The known hazards related to chemicals used in the lab are found in:(a) PCM - Poison Control Manual (b) MSDS - Material Safety Data Sheets(c) Chemical Safety Data Sheets

9. What should be done if a chemical gets in your eye?(a) Rinse your eyes with water from the eyewash fountain for at least 15 minutes(b) Rinse under the safety shower for 5 minutes(c) Nothing, unless the chemical causes discomfort

10. Why should contact lenses never be worn in the lab?(a) They could inadvertently fall out of the eye(b) They are too hard to find if they fall out (c) Chemical vapour could react with or become trapped between the eye and the lens

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11. Why should you wear goggles in the lab even if you personally are not working with anychemicals?(a) Discomfort is a part of science (b) They make you look smart(c) Someone at another workstation may be working with chemicals and splash you

12. Broken glass and sharp objects must be disposed of:(a) In the trash can (b) In the yellow glass container at the back of the lab(c) In the back of someone else’s locker when they are not looking

13. What precautions are needed with long hair and beards?(a) Must be shampooed (b) No long hair and beards allowed in the lab(c) Keep long hair tied back and away from flames and chemicals

14. To dispose of waste chemicals:(a) Pour them down the sink & flush with lots of water(b) Hide them in the back of your locker(c) Put them in the designated waste containers provided

15. When heating with a burner or hot plate in the lab, NEVER:(a) Put hot glass on a cold counter top (b) Leave it unattended(c) Heat a closed container (d) a, b and c

19. If you spill a small amount of chemical on your skin or clothing you should:(a) Evacuate the lab immediately (b) Take a shower when you get home(c) Flush with water for 5 minutes and notify your TA immediately

20. If a large chemical spill occurs on your skin or clothing you should:(a) Remove the clothing, rinse under the safety shower and notify your TA(b) Run madly about the lab until it dries (c) Call 911

21. You must report any form of personal injury incurred in the lab to your TA.(a) True (b) False

22. The person ultimately responsible for your safety in the lab is:(a) The safety committee (b) Your lab partner (c) Your TA (d) Yourself

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POLICY ON PLAGIARISM

Below are the guidelines that outline the policy on plagiarism for the University of Windsor. TA’swill address this issue with students during the first week of the laboratories and discuss theoutcomes that are of a result of plagiarism.

THE COPYING OF WEB PAGES IS CONSIDERED PLAGIARISM AND IS NOT ACCEPTABLE.

ACADEMIC INFORMATIONUndergraduate Degree Regulations

2.4.22 POLICY ON PLAGIARISM

Plagiarism is defined as: "The act of appropriating the literary composition of another, orparts of passages of his or her writing, or the ideas or language of the same, and passingthem off as the products of one's own mind." (Black's Law Dictionary).

It is expected that all students will be evaluated and graded on their individual merit andall work submitted for evaluation should clearly indicate that it is the student's own contribution.

Students often have to use the ideas of others as expressed in written or published workin preparing essays, papers, reports, theses and publications. It is imperative that both thedata and ideas obtained from any and all published or unpublished material be properlyacknowledged and their sources disclosed. Failure to follow this practice constitutesplagiarism and is considered to be a serious offence. Thus, anyone who knowingly orrecklessly uses the work of another person and creates an impression that it is his or herown, is guilty of plagiarism.

Plagiarism also includes submitting one's own essay, paper, or thesis on more than oneoccasion. Accordingly, it is expected that a thesis, essay, paper or a report has not beenand is not concurrently being submitted for credit for any other course. In exceptionalcircumstances and with the prior agreement of the instructor, a student may use researchcompleted for one course as part of his or her written work for a second course.

A confirmed incident of plagiarism will result in a sanction ranging from a verbal warning,to a loss of credit in the course, to expulsion.

Source: 2004/2006 Undergraduate Academic Calendar, Paragraph 2.4.22

For additional information on this topic, please visit the University of Windsor’s websitededicated to student integrity at http://www.uwindsor.ca/aio

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SENATE BY-LAW 31

UNIVERSITY POLICY IN RESPECT TO JUDICIAL PROCEDURESARTICLE I. SANCTIONS AND DEFINITIONS

Proscriptions StatedUniversity discipline is limited to student misconduct which adversely affects the Universitycommunity's pursuit of its educational objectives. Students are expected to conductthemselves in a manner compatible with the objectives and purposes of the University ofWindsor. Any student at the University of Windsor whose conduct is improper in that itexhibits a lack of integrity touching upon the educational objectives and requirements ofthe University must be disciplined appropriately in the interest of safeguarding andupholding these standards.

It is desirable to define and identify further the standards demanded of each student at theUniversity of Windsor in the interest of educational integrity. Enumerated below areillustrations of improper conduct which would lead to an inference of lack of integrity. Theseare illustrative only and shall not be taken as in any way limiting the generality of the highstandards of conduct required by the objectives and purposes of the University of Windsor.

Examples of misconduct for which students are subject to university discipline are definedas follows:

1. Dishonesty, such as cheating, plagiarism, impersonation at an examination, orknowingly furnishing false information to the University.

2. Forgery, alteration, or use of University documents, records, or instruments ofidentification with intent to defraud.

3. Intentional obstruction or disruption of teaching, research, administration,disciplinary proceedings, or other University activities, including public servicefunctions, and other authorized activities on University premises.

4. Malicious abuse of any person on University premises or at University sponsoredor University supervised functions or malicious conduct which threatens, endangersor harasses any such person.

5. Theft from or deliberate damage to University premises or theft of or deliberatedamage to property of a member of the University community on Universitypremises.

6. Failure to comply with directions of members of the University administration or ofthe teaching staff acting in the proper performance of their particular duties.

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7. Violation of published University regulations, including regulations relating to entryand use of University facilities.

8. Violation of published rules governing University residence halls.

9. Deliberate alteration or misappropriation of computer records, data, software, etc.of the University or member of the University community.

10. Breach or misuse of the Code of Computing Practice for the University of WindsorComputer Centre user.

Sanctions Defined

1. Admonition. Notice to the student, orally or in writing, that s/he has violatedstudent rules and that continuation or repetition of the conduct found wrongful,within a period of time stated in the warning, may be cause for more severedisciplinary action.

2. Censure. Written reprimand for violation of a specified regulation, including thepossibility of more severe disciplinary sanction in the event of conviction for theviolation of any University regulation within a period of time stated in the letter ofreprimand.

3. Disciplinary Probation. Exclusion from participation in privileges or extracurricularUniversity activities as set forth in the notice of disciplinary probation for a specifiedperiod of time.

4. Restitution. Reimbursement for damage or misappropriation of property.Reimbursement may take the form of appropriate service to repair or otherwisecompensate for damages.

5. Suspension. Exclusion from classes and other privileges or activities as set forthin the notice of suspension for a definite period of time.

6. Expulsion. Termination of student status for an indefinite period. The conditionsof readmission, if any is permitted, shall be stated in the order of expulsion.

7. Exclusion from Campus. Denial of access to the campus for an indefinite periodfor non-academic misconduct. The conditions for removing this ban, if any, shall beincluded in the exclusion order.

Source: http://athena.uwindsor.ca/units/senate/senate.nsf/Bylaws?OpenView

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STUDENT CONTRACT

I, ___________________________________, have read and understood the

Laboratory Safety and Plagiarism Policy sections of this manual and agree to abide

by the dictates of these documents. I realize that failure to do so may result in my

dismissal from the lab with no opportunity to make-up missed work. I, also, understand

the serious consequences that are to be taken if I have been caught with plagiarism

and I will take full responsibility for my actions. In addition, I have successfully

completed the Safety Quiz.

* * * * * * * * * * * * * * * * * * * * * * * * * * * * * * * * * * * * * *

Course: 59-_________ Section: ________ Date: _______________________

TA’s Name: __________________________________________________________

Student’s Name (please print): __________________________________________

ID # ___________________ Signature ___________________________________

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MARKING SCHEME AND OUTLINEAll labs are marked out of 20.Abstract (3 marks)

The abstract is a crucial part of your laboratory report. It is a clear written paragraphwhich briefly conveys the intent of the laboratory exercise, mentions the methodology utilizedduring the laboratory period, summarizes important experimental observations, findings andconclusions. The abstract length must be between 200 to 300 words.

Procedure (0 marks)No marks are given for this section. You may indicate “as outlined in the 59–240 lab

manual”, quoting the relevant pages. However, if there is a major deviation from experimentalprocedure or significant problem with the data arising from experimental error, you must statethese differences in this section, as these deviations could influence the marks you receive in latersections (i.e., see Results and Calculations below).

Results and Calculations (Labs 1-4, 8 marks & Lab 5, 6 marks)Present your data in neatly laid out tables. Always include units and estimated errors

arising from experimental errors, instrumental readings, etc. Include a comparison of yourexperimental data with literature values where appropriate. Show a sample calculation for eachtype of calculation required, including the equation used, the numbers substituted into theequation, unit analysis and the final answer.

Discussion (Labs 1-4, 7 marks & Lab 5, 9 marks)This section is meant to convey your understanding of the experiment. Discuss all

important results, where trends and anomalies are discussed with reference to the underlyingtheory of the experiment. Important results should be compared with those found in theliterature, and reasons for differences between experimental and literature data should bediscussed. When there are several sources of possible error, indicate to what extent each mayhave affected the final results. You may also choose to include in your discussion possibleapplications of the experiment, suggested improvements or extensions to the experiment, and/oralternative methods of measurement. Answer the lab questions at the end of each experiment!

Conclusions (1 mark)Write a brief statement summarizing the most important results of the experiment,

including numerical final results and per cent errors (if applicable).

References (1 mark)Make a list of all reference material, using the same format as the lab manual. References

should appear in the order they appear in the text. References in the text should appear as super-scripted arabic numerals. For example: The boiling point of benzene is 80.1 oC.3

Reference format: 1. Atkins, P.W. Physical Chemistry, 7th Edition, Oxford University Press, Oxford, 2001.

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GENERAL LABORATORY OUTLINE1. Lab attendance is compulsory.2. Laboratory reports must be turned in at the start of the next lab period. In the final week,

reports must be handed in exactly one week after the completion of the lab.3. Each laboratory report is marked out of 20, and accounts for a maximum possible 3

marks towards your final grade (the five laboratories have a total weight of 15 % in thiscourse).

4. Laboratories which are not completed receive a grade of zero. A completed laboratoryrefers to both attendance and performance in the lab, as well as handing in a completedreport.

5. If the laboratory is not completed due to illness or family tragedy, the student may presenta doctor’s note or some other documentation to the upper year lab coordinator in order tobe excused from the laboratory. In this case, the student will receive a mark on this“excused laboratory” equal to the average mark on all other laboratory reports.

6. No make-up labs are offered.7. If there is any evidence of plagiarism or duplication in lab reports among lab partners,

friends or other students, all parties will receive a grade of zero and be subject toacademic discipline, as per the University of Windsor Undergraduate Calendar.

Prelab PreparationRead your laboratory manual carefully. You must be prepared to answer questions asked

by your TA about the experiment and for surprise quizzes. You must know, or have writtendown, the physical constants of the chemicals you will be using. MSDS sheets are readilyavailable for those who wish to consult them. You must also know how to prepare solutionsrelevant to your experiment (e.g. 4M HCl from concentrated HCl) before you come to the lab. Ifthere is any procedure, or equipment operation, with which you are not familiar, PLEASE ASKYOUR TA FOR HELP. This will help to prevent accidents and waste of valuable time.

DISPUTES, COMPLAINTS, AND POLICYAny disputes or complaints arising from the laboratory course should, in the first

instance, be drawn to the attention of the TA. Should his/her decision be deemed unjust,representations may be made, in turn, to the Professor instructing the course, the Head of theDepartment of Chemistry and Biochemistry, and the Dean of Students.

ACKNOWLEDGEMENTSThe current edition of this manual was prepared by Robert W. Schurko. Experiment 5

was developed by Mohammad Harati. Previous editions of this manual, as well as developmentand refinement of the experiments are credited to Cory Widdifield and Robert W. Schurko.

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dpdT

'∆trs HT∆trs V [1]

ln XA '∆HA

R1

TA

&1T [2]

ln XB '∆HB

R1

TB

&1T [3]

XA % XB ' 1 [4]

EXPERIMENT 1SOLID – LIQUID PHASE DIAGRAM

Important: bring a formatted 3.5” floppy diskette/USB flash drive for this laboratory – you willneed it to save your data files!

IntroductionThe relation of cooling curves to phase diagrams form the basis of “thermal analysis”, an

important technique for constructing phase diagrams. In the solid-liquid phase equilibriumchosen for study here, the two components, although miscible with one another in the liquidphase, are of limited solubility in one another as solids. Thus, we can consider them as puresolids, plus a two-component liquid. Such systems exhibit an eutectic temperature at which thethree phases can coexist in equilibrium at a fixed pressure.

Solid-liquid equilibria differ from liquid-vapour equilibria in that they are essentiallyindependent of pressure changes on the order of a few atmospheres, owing to the small molarvolume change associated with fusion. This is a consequence of the Clapeyron equation:

We shall be concerned here with temperature-composition diagrams at p = 1 atm.If the liquid solution behaves ideally, the “solubility” of each component in the liquid

depends on temperature, according to:

where:XA and XB are the mole fractions of components A and B, respectively,∆HA and ∆HB are the heats of fusion of components A and B, respectively, andTA and TB are the melting points of the pure components.

Equations [1] and [2] can be represented in the following phase diagram:

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Figure 1.1

The eutectic composition (XE) and eutectic temperature (TE) are given by the intersectionof the two liquid curves. With the aid of graph (a) in Figure 1.1, we can predict the generalnature of the cooling curves in a system of this kind. The curves in graph (b) are plots oftemperature against time obtained when liquid solutions of various compositions are allowed tocool.

When a liquid consisting of pure A is cooled, the temperature falls until solid A begins toform. The temperature then remains constant until solidification is complete, where upon it fallsagain. It is said that the curve shows a thermal arrest. When a liquid having the eutecticcomposition is cooled, the behaviour is similar in that a thermal arrest is obtained.

However, when a liquid of some other composition – for example, X1 (see graph (a)) iscooled, solid A begins to form at temperature T1. This tends to deplete the liquid of componentA, so that its composition passes through X2, X3, ... and the temperature falls as long as solid Aalone continues to come out of solution.

The abrupt change in slope, which occurs when solid A begins to form, is called a break. When the composition of the solution finally reaches XE, solid B begins to form together withsolid A and the two solids continue to separate from solution at the temperature TE until no liquidremains, and thus an arrest occurs.

The binary solid-liquid phase diagram for the naphthalene-diphenylamine system will beconstructed from cooling curves. Several mixtures of different ratios of the two components willbe melted, and temperature versus time curves will be plotted as the mixtures cool. Thetemperatures at which these breaks and arrests occur are plotted as a function of composition ofthe mixtures to obtain the phase diagram.

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Materials

• 1 – 600 mL beaker• 1 – 200 mL beaker• 1 – plexiglass container and lid• 1 – large test tube• 1 – medium test tube• 1 – jumbo magnetic stirrer• 1 – magnetic stirring bar

• 1 – hot plate• 1 – rubber sleeve• boiling chips• Vernier with temperature probe• naphthalene (C10H8)• diphenylamine (C12H11N)

ProcedureThe apparatus will be set up as shown in Figure 1.2, and time-temperature curves should

be measured for the mixtures outlined in Table 1.1. Follow the procedure outlined below toconstruct these curves. Instruction on the use of the Logger-Pro software will be provided byyour TA.

Table 1.1

Run diphenylamine (g) naphthalene (g) made by

1 0 5

2 1 5 run (1) + 1.0 g diphenylamine

3 2.5 5 run (2) + 1.5 g diphenylamine

4 5 5 run (3) + 2.5 g diphenylamine

5 10 5 run (4) + 5.0 g diphenylamine

6 5 0

7 5 1 run (6) + 1.0 g naphthalene

8 5 1.67 run (7) + 0.67 g naphthalene

1. Fill the 600 mL beaker 3/4 full with water and bring to a boil. Be sure to use boilingchips to prevent bumping.

2. Fill the plexiglass container with cold water and maintain the temperature between 10-20EC throughout the course of the lab.

3. Weigh all samples to an accuracy of 0.01 g.4. Transfer the first sample into the inner test tube. Insert the temperature probe and

magnetic stirring bar into the inner test tube as well.5. Place the test tube into a hot water bath and heat until the solid is completely melted.6. Dry the test tube and insert through the rubber sleeve of the large test tube contained

within the cold water bath in the plexiglass container (Figure 1.2). Turn on the jumbomagnetic stirrer to ensure that the sample is continually stirring.

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jumbo magnetic stirrer

ice water bath

to Labpro

temperature probe

stir bar

plexiglasscontainerand lid

medium test tube

600 mL beaker

hotplate

Figure 1.2

7. Click on on the Logger Pro software. This will provide a real time plot of thechanging temperature in the sample. The collection will continue for a maximum of 1000seconds. However, if you observe a distinct arrest or break in the pattern, you mayterminate the acquisition and begin a new run. Note: keep the water bath cold, especiallyfor the later runs! Remember that the pure compounds and the eutectic mixture shouldhave only one arrest! Be sure to pull out the stir bar before dumping waste!

8. Save a file for this run by exporting the data as a text file. Excel can later be used toprocess this data. Make sure to use file names that will clearly distinguish your data sets.

9. Find the arrests on the curves for the pure compounds, the breaks on all mixtures, and theeutectic arrest for two of the latter.

Calculations1. Extract the break and arrest temperatures from the cooling curves. Print out plots of your

cooling curves (Excel can be used). Convert all temperatures to the thermodynamictemperature scale (i.e., Kelvin) and use these units for all of the calculations.

2. Determine the mole fraction, Xnaphthalene, for all of the mixtures. Determine the molefraction of diphenylamine, Xdiphenylamine, for runs 6, 7 and 8.

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3. Using the break and arrest temperatures of these mixtures, plot the two “limbs” of thesolid-liquid phase diagram (i.e., temperature as a function of the mole fraction ofnaphthalene). Runs 1-5 will constitute one “limb”, and runs 6-8 will make up the other.

4. Draw the liquidus curves, identify the eutectic line and identify the phases present in eacharea of the diagram.

5. Determine the eutectic composition and eutectic temperature from the phase diagram.6. Using [1] and [2], plot ln Xnaphthalene against 1/T for runs 1-5 and plot ln Xdiphenylamine against

1/T for runs 6-8. The plots should be linear with:slope = -∆Hi/R and-slope/intercept = Ti (the melting temperature of the pure component.)

7. Calculate the enthalpy of fusion and the melting point for each component, assuming thatan ideal liquid solution is formed. Compare to literature values by calculating the percenterror for ∆fusH and temperature, as well as the absolute error for the temperature.

Lab Questions1. A series of Ni-Mn mixtures were prepared and allowed to reach equilibrium at various

temperatures. Use the following data to plot a phase diagram (preferably on an Excelspreadsheet) for the Ni-Mn system. Label the different components and number ofphases in each part of the phase diagram.

Mn-Rich mixtures:

T/(oC) 1260 1200 1150 1100 1050 1000

w(Ni)sol 0.00 0.04 0.08 0.13 0.22 0.45

w(Ni)liq 0.00 0.07 0.12 0.18 0.29 0.45

Ni-Rich mixtures:

T/(oC) 1050 1100 1150 1200 1250 1300 1350 1400 1450

w(Ni)sol 0.58 0.64 0.70 0.75 0.80 0.85 0.90 0.96 1.00

w(Ni)liq 0.54 0.62 0.68 0.73 0.78 0.83 0.88 0.94 1.00

2. What is the composition of the solid solution in equilibrium with a liquid mixture of Niand Mn at 1000 oC? What is this point designated as?

References9. Atkins, Peter and Julio de Paula. Physical Chemistry. 7th ed. New York: W. H.

Freeman, 2002. 144-148.10. Shoemaker, David P., Garland, Carl W., and Joseph W. Nibler. Experiments in Physical

Chemistry. 6th ed. New York: McGraw–Hill, 1996. 215-221.11. Silbey, Robert J., and Robert A. Alberty. Physical Chemistry. 3rd ed. Wiley, 2001.

Chapter 6, Section 6.9.

20

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21

[ ][ ][ ]Ka

3

3

H CH3COO

(CH3COOH)

H CH COO

CH COOH= ×

+ −+ −γ γ

γ( ) ( ) [1]

EXPERIMENT 2DETERMINATION OF Ka USING THE CONDUCTANCE METHOD

IntroductionEquilibrium Processes

When a pure sample of liquid-state acetic acid (i.e., CH3COOH(l)/HAc(l)) is added to abeaker of pure water, at least two significant processes occur. First, the acetic acid (assuming itto be the solute) will dissolve completely in the water (assuming it to be the solvent). As aceticacid is infinitely soluble in water (i.e., the two liquids are miscible), this process of dissolutionwill result in a solution of acetic acid and water with no measurable amount of pure solute. Asthis process has proceeded completely, (i.e., 100 %) an equilibrium expression need not beconsidered. Second, the dissolved acetic acid will undergo the process of dissociation, a processwhich may be represented by an equilibrium expression and which is the focus of this lab. Bothprocesses may be expressed symbolically:

When acetic acid dissociates, a proton is liberated, and therefore the equilibrium constantrepresenting this process is referred to as the acid dissociation constant (Ka).

where the parameter γ is defined as the activity coefficient of that particular substance.

ConductanceA very common method to use when determining the acid dissociation constant of a

sample is called the conductance method, as one measures the conductance of a solution. Conductance, G, (SI unit is the seimens, S, where 1 S / 1 Ω-1) is defined as the reciprocal ofresistance and can be understood to represent the ease with which electrical current flows througha given substance. With regards to our setup, the conductance is seen to be proportional to theelectrode surface area, A, and inversely proportional to the distance between the electrodes, l:

where κ is the conductivity (µS cm-1) of the solution. The conductivity parameter depends upon(i) the equivalent concentrations, (ii) the charge numbers and (iii) the mobilities of the ionicspecies present. An additional parameter, α, representing the fractional ionization at a givensolute concentration is now introduced and is related to conductivity:

where:

CH3COOH(l)

H2O(l)CH3COOH(aq) CH3COO-

(aq) H+(aq)

G ' κ Al [2]

κ ' 1000Cανö (U%% U

&) [3]

22

# C is the concentration of the solute in solution. (mol L-1)# v is the simplified form of a parameter which represents the number of equivalents

and the charge numbers of positive and negative ions per mole of electrolyte andis equal to 1 for any 1:1 electrolyte.

# ö is the Faraday constant. (C equiv-1)# U± are the ionic mobilities for the positive and negative ions (in m2 s-1 V-1)

By rearranging [3], one obtains the equivalent conductance, Λ (Ω-1 cm2 equiv-1). Thisform is useful as it eliminates the need to worry about ionic mobilities and charges. Please notethat this is a simplification and may be applied only to 1:1 electrolytes. The equivalentconductance is defined as:

For this system, it is assumed that the fractional ionization is approximated by the ratio ofequivalent conductance at a given concentration to that at infinite dilution:

For strong electrolytes, Λ varies only slightly with concentration (as α does not vary), andΛ0 can be determined by measuring Λ at various solute concentrations:

where ε is a proportionality constant. Thus, a plot of Λ against , followed by extrapolation toCinfinite dilution, will produce Λ0. For weak electrolytes, Λ is no longer nearly constant (as αvaries markedly with concentration) and Λ0 must be determined using an indirect approach,which relies upon the equivalent conductances of several strong electrolytes at infinite dilution(The Kohlrausch method):

Relating Conductance Measurements to Dissociation ConstantsRecalling that C represents the concentration of the undissociated acetic acid and α

represents the fractional dissociation of acetic acid, it is understood that:

This modifies [1] in the following manner:

Λ 'κ

1000C' αö (U

%% U

&) [4]

α – ΛΛ0

[5]

Λ ' Λ0(1 & ε C) [6]

Λ0(HAc) ' Λ0(HCl) % Λ0(NaAc) & Λ0(NaCl) [7]

[H %] ' [CH3COO &] ' αC [8]

Ka 'α2C

(1 & α)×

γH %γCH3COO &

γCH3COOH[9]

23

Recalling the definition of fractional dissociation in terms of equivalent conductances, the aboveequation modifies to:

Let the right hand portion of [10] which does not contain activity coefficients equal thedissociation function, Kc, which is a quantity related to the true acid dissociation constant. Asundissociated acetic acid is necessarily non–electrolytic, its activity coefficient equals unity andmay be disregarded. Also, as all solutions are of relatively low solute concentration, one may usethe Debye-Hückel limiting law to give:

By determining the dissociation function for a number of acetic acid solutions of lowconcentration and by plotting log(Kc) versus with linear extrapolation to C = 0, one canCαdetermine the acid dissociation constant, Ka, at that particular temperature.

The Dissociation Constant as a Function of TemperatureThe acid dissociation constant may be treated just like any other equilibrium constant and

thus its variation with changing temperature obeys the van't Hoff equation, which takes the formbelow after it is integrated:

It is assumed that the enthalpy of dissociation, ∆dH, is constant over the temperaturerange, [T1, T2].

Ka 'Λ2C

Λ0(Λ0 & Λ)×

γH %γCH3COO &

γCH3COOH[10]

log Ka ' logKc & 1.018 Cα [11]

ln Ka2& ln Ka1

' &∆dH

R1T2

&1T1

[11]

24

Materials

• 3 – 250 mL volumetric flasks• 1 – 100 mL graduated cylinder• 1 – 50 mL graduated cylinder• 5 – 100 mL beakers• 1 – 250 mL beaker

• Vernier with conductance probe• hydrochloric acid (HCl)• acetic acid (CH3COOH)• sodium acetate (CH3COONa)• sodium chloride (NaCl)

ProcedureThe following stock solutions will be provided:1. 0.1 M sodium chloride (NaCl)2. 0.1 M sodium acetate (CH3COONa)3. 0.04 M hydrochloric acid (HCl)4. 0.02 M acetic acid (CH3COOH)

A. Preparation of Required Solutions1. Each stock solution provided will be used to make three dilute solutions, each 250 mL in

volume. In addition, keep about 50 mL of the stock solution in a labelled 100 mL beaker. All required solution concentrations are outlined in Table 2.1 (Note: prepare only one setof solutions at a time - only use deionized water, do not use tap water).

Table 2.1 - Required Solution Concentrations

Set Compound Concentrations (M)

1 NaCl 0.02, 0.04, 0.08, 0.10

2 CH3COONa 0.02, 0.04, 0.08, 0.10

3 HCl 0.005, 0.01, 0.02, 0.04

4 CH3COOH 0.004, 0.008, 0.01, 0.02

B. Calibration of the Conductivity Probe (repeat steps as required)1. Make sure the computer console has been turned on and that the conductivity probe has

been correctly connected. Please ask a TA for help if this has not yet been done.2. Load the Logger Pro software and make sure that the conductivity probe is set to the 0 –

20000 µS range if using a strong electrolyte or the 0 – 2000 µS range if using a weakelectrolyte. The switch controlling the range is found on the black conductivity probebox.

3. Rinse the probe with de-ionized water (do not scratch the electrodes!) and dry theoutside with a delicate task wiper (i.e., a Kimwipe).

4. Load an appropriate calibration file, which can be found under the “Experiments/Probes& Sensors/Conductivity” directory. Note: consider the conductivity range you are using.

25

This calibration file will automatically setup the probe for use. Make note of thisconductivity reading (“the air reading”) as it must be subtracted from all subsequentmeasurements.

5. As a quick check of the probe, (perform only once) add about 50 mL of de-ionized waterto a 100 mL beaker and about 50 mL of tap water to another 100 mL beaker. Take theconductivity measurements of both solutions. The de-ionized water reading should beequal or very close to the air reading, while the tap water reading should be muchdifferent. Also, using the Vernier temperature probe, measure and then record thetemperature of one of the solutions into your data sheet.

C. Conductivity Measurements1. Once a set of solutions has been prepared and making sure that the probe has been

correctly calibrated, add about 50 mL of each diluted solution to a labelled 100 mLbeaker. A 250 mL beaker has been provided to use as a quick rinse and should beprepared with fresh de-ionized water for each set of solutions. For acetic acid: the de-ionized water should be changed after each measurement.

2. Dip the conductivity probe into the solution (see Figure 2.1) with the lowest electrolyteconcentration and click on the button, which can be found near the top rightof the screen. This will start the data acquisition process and produce a plot on thescreen.

3. Once the reading has stabilized (ca. 10 s), record the value on your data sheet, remove theprobe from the solution, blot dry, swirl in the de-ionized water bath, blot dry again anddip the probe into the solution having the next highest electrolyte concentration.

4. Repeat step 3 until the conductivities of all the solutions in a set have been recorded.5. Clean all glassware (beakers, flasks, etc.) and move on to the next set of solutions.6. After the conductivities of all 16 solutions (4 sets of 4 solutions) have been measured,

perform a final cleaning of the conductivity probe, exit the Logger Pro program andreturn all glassware to the appropriate student locker.

Figure 2.1

26

Calculations1. Using equation [4], determine Λ values for all of the solutions that contain a strong

electrolyte. The cell constant, A/l, is equal to 1 cm.2. Plot Λ vs. and determine Λ0 for all the strong electrolytes.C3. Using the Kohlrausch method, evaluate Λ0 for acetic acid.4. For each acetic acid solution, determine Λ, α, and Kc, and present these findings in a

table.5. By plotting log(Kc) vs. and extrapolating to C = 0, find pKa for the temperature atCα

which the experiment was conducted.6. Determine the enthalpy of dissociation, ∆dH, for acetic acid under standard conditions.7. Using [12], determine Ka under standard conditions and compare your findings with the

literature. Note: the value of Ka for acetic acid is incorrect in both Atkins 7th and 8th

editions, though the listed pKa is correct.

Lab Questions 1. Why, when determining Λ0 for a weak electrolyte, it is preferable to determine the

equivalent conductance at infinite dilution for several strong electrolytes as opposed todetermining it directly using the weak electrolyte?

2. Using the Debye-Hückel-Onsager limiting law, compute the ionic molar conductivity forZn2+(aq) and SO4

2-(aq) ions in a 0.01 molar aqueous solution of ZnSO4 at 298.15K. Thelimiting ionic molar conductivities are λ(Zn2+) = 105.6 S cm2 mol-1 and λ(SO4

2-) = 160.0 Scm2 mol-1. Use your results to calculate Λ(ZnSO4) at the given concentration. Calculatethe percent error in your result by comparison to the result in the CRC Handbook.

3. You should have noticed that, at the end of the experiment, only a small fraction of thetotal solution made was actually utilized for conductance measurements. Why were suchlarge amounts of solution prepared when only a small portion was used?

4. When the sample under study is a non 1:1 weak electrolyte, explain why the definition ofequivalent conductance, as outlined earlier, no longer holds.

5. The molar conductivity of an aqueous 0.10 molar solution of AgNO3 is 109.09 S cm2 mol-

1 at 298.15K. When this solution is placed in a particular conductance cell, the resistanceof the solution is found to be 35 Ω. Compute the specific conductivity of the AgNO3solution.

References1. Weast, Robert C. “Physical Constants of Organic Compounds.” CRC Handbook of

Chemistry and Physics. 56th ed. 1976. C-76.2. Shoemaker, David P., Garland, Carl W., and Nibler, Joseph W. Experiments in Physical

Chemistry. 6th ed. New York: McGraw–Hill, 1996. 228-238.3. Atkins, Peter and Julio de Paula. Physical Chemistry. 7th ed. New York: W. H.

Freeman, 2002. 235-238, 833-838.

27

Dat

a Sh

eet

Tem

pera

ture

of s

olut

ions

: ___

____

_ EC

Con

cent

ratio

n of

CH

3CO

OH

solu

tion:

___

____

_ m

ol L

-1

Con

cent

ratio

n of

HC

l sol

utio

n: _

____

___

mol

L-1

Con

duct

ivity

read

ing

of a

ir us

ing:

(i)

0 –

200

0 µS

cm

-1 ra

nge

____

____

____

µS

cm-1

(ii

) 0 –

200

00 µ

S cm

-1 ra

nge

____

____

____

µS

cm-1

#1 -

NaC

l#2

- C

H3C

OO

Na

#3 -

HC

l#4

- C

H3C

OO

H

Con

c. (M

)κ

(µS

cm-1)

Con

c. (M

)κ

(µS

cm-1)

Con

c. (M

)κ

(µS

cm-1)

Con

c. (M

)κ

(µS

cm-1)

0.02

0.02

0.00

50.

004

0.04

0.04

0.01

0.00

8

0.08

0.08

0.02

0.01

0.1

0.1

0.04

0.02

28

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29

2A % 4B ! 3C % D [1]

rate ' ν 'd[D]

d t'

13

d[C]d t

' – 14

d[B]d t

' – 12

d[A]d t

ν ' k [A]a [B]b [2]

EXPERIMENT 3THE IODINE CLOCK

IntroductionThe Rates of Chemical Reactions

Broadly defined, chemical kinetics is the study of the rates at which chemical reactions proceed. Oftentimes, reaction rate data helps chemists to develop reaction mechanisms for agiven chemical reaction. The rate of a chemical reaction depends upon several parameters, suchas temperature, pressure and the initial composition of a system undergoing a reaction. If onewishes to consider a reaction for which a balanced chemical equation is known, several reactionrates can be determined:

In [1], it is observed that 2A*s and 4B*s are required to produce 3C*s and one D. Fordilute solutions, reaction rates can be expressed by considering the change in concentration of aparticular constituent that is being consumed or produced with respect to the change in time:

Reaction Rate Law and Reaction OrderThe rate, ν, at which a chemical reaction proceeds is often found to be proportional in

some fashion to the concentration of each reactant raised to some exponential power:

Where a and b are experimentally determined values and k is the rate constant.

The rate law is not, in general, to be inferred from a balanced chemical reaction, rather, itmust be determined experimentally. The constants a and b are oftentimes integers and theoverall reaction order is equal to their sum. Thus, for example, if a and b were determined to beequal to 2 and 1, respectively, the reaction is said to be third-order (a + b = 2 + 1 = 3).

Determination of the Rate LawOne very popular method of determining the rate law is known as the method of initial

rates, and is aptly titled as the rate is measured only for the initial stages of a chemical reaction. This method involves making a series of solutions, each with a different composition andobserving the initial reaction rate as a function of the changing reactant composition. For areaction which involves two reactive species, the general process involves first keeping theconcentration of one of the reactants constant (reactant A, for example) while varying theconcentration of the other (reactant B). Then, the concentration of B is held fixed while varyingthe concentration of A. By plotting the logarithm of the initial rate, ν0, against the logarithm ofthe concentration of the reactant whose concentration is being varied, the reaction order withrespect to that constituent can be determined as it is equal to the slope of a first-order line of bestfit (m = a, according to [2], for the case where [A] is being varied; see Figure 3.1).

30

S2 O 2–8 (aq)

% 2I –(aq)

! 2SO 2–4 (aq)

% I2 (aq) [3]

log(v0) = 2.0005log([A]) - 1.8429R2 = 1.00

-6.5

-6.2

-5.9

-5.6

-5.3

-5

-4.7

-2.4 -2.2 -2 -1.8 -1.6 -1.4

log [A] (units of M)

log

v 0 (un

its o

f M s

-1)

Figure 3.1 Plot of log ν0 versus log [A]. According to the line of best fit, the reaction is clearlysecond-order with respect to reactant A.

I2 (aq)% 2S2 O 2–

3 (aq)! S4 O 2–

6 (aq)% 2I –

(aq) [4]

The Iodine ClockIn this lab experiment, the reaction between the persulfate ion (S2O8

2–) and the iodide ion(I –) is studied in aqueous media and may be represented by the following chemical equation:

Two additional reagents are used: the first provides the thiosulfate anion (S2O3 2–). This

anion reacts very quickly, with respect to the process described by [3], with any iodine present:

Once all of the thiosulfate is consumed, the iodine being evolved at this point will reactwith any iodide present and their product forms a complex with the second additional reagent,which is starch. This reaction produces a purple colour which signals the end of the datasampling period. Therefore, by considering the amount of thiosulfate added to the reactionmixture, one can measure the rate of the reaction.

31

ln k ' ln A –Ea

RT[6]

k ' κ kTh

RTpEG

e– ∆‡G

RT ' κ kTh

RTpEG

e–∆‡H

RT e∆‡SR [7]

Ea ' ∆‡ H % RT [8]

logk ' logk0 % 1.018 zAzB I [5]

The Reaction MediumOne should also consider the influence that the reaction medium has upon the reaction

rate. The observed reaction rate is not only a function of the items outlined earlier, but is alsodependent upon the solvent used (not considered in this experiment) and the ionic strength, I, ofthe solution. In order to determine the rate constant at infinite dilution, k0, a correction whichaccounts for the ionic strength of the solution is done (note that this form of the correction is onlyapplicable when the Debye-Hückel limiting law holds):

Where zA and zB are the charge numbers associated with the two atoms or molecules that areundergoing the chemical reaction under study.

The Rate Constant as a Function of TemperatureThe rate constant for many chemical reactions is found to increase in a linear fashion

when a plot of ln k versus 1/T is composed. This is the graphical expression of the Arrheniusequation:

where A is commonly denoted as the frequency factor and where Ea is the activation energy ofthe reaction under consideration.

Activated Complex TheoryIn bimolecular reactions, the two reacting species (A and B, for example) are postulated

to pass through some form of transition state, which is often higher in energy than the reactants. This is the basis of activated complex theory (sometimes referred to as transition state theory). By using this theory, one can relate the rate constant to the Gibbs energy of activation, ∆‡G, for aparticular temperature:

Where κ is the transmission coefficient, a dimensionless quantity which will be assumed to equal1, k is the Boltzmann constant, h is the Planck constant, and ∆‡H and ∆‡S are the enthalpy ofactivation and the entropy of activation, respectively. All other parameters carry their usualmeaning. For reactions carried out in solution, the enthalpy of activation is related to theactivation energy:

Therefore, with ∆‡G and ∆‡H known, ∆‡S can be determined by using the equation whichexpresses the change in Gibbs energy for an isothermal process.

32

Materials

• 4 – 50 mL burettes• 3 – medium test tubes • 1 – test tube rack• 6 – 100 mL beakers• 1 – 1 L beaker• 3 – 100 mL volumetric flasks• 3 – 100 mL Erlenmeyer flasks• 1 – hot plate & stirrer• 1 – water bath • 1 – thermometer• 2 – scoopula

• 2 – magnetic stir bars• 1 – stopwatch• 1 – 1.0 mL pipette & bulb• 2 – 5.0 mL pipettes• ammonium chloride (NH4Cl)• potassium chloride (KCl)• ammonium persulfate ((NH4)2S2O8)• potassium iodide (KI)• sodium thiosulfate pentahydrate

(Na2S2O3C5H2O )• starch solution

ProcedureThe following solutions will be provided:1. 0.15 M ammonium chloride (NH4Cl) (A1)2. 0.12 M potassium chloride (KCl) (A2)

A. Preparation of Required SolutionsPrepare the following solutions using the room temperature de-ionized water (note: DO NOTuse tap water for the preparation of these samples):1. 100 mL of 0.05 M ammonium persulfate ((NH4)2S2O8) (A3)2. 100 mL of 0.12 M potassium iodide (KI) (A4)3. 100 mL of ca. 0.007 M sodium thiosulfate pentahydrate (Na2S2O3C5H2O) (A5)

33

B. The Effect of Reactant Concentrations on ν01. For A3 and A4, transfer some of each solution into separate 100 mL beakers and then into

separate 50 mL burettes. You may wish to label the burettes so you do not mix up thesolutions.

2. In a similar fashion as step B1, transfer some de-ionized water, as well as some starchsolution, into 50 mL burettes. A total of four burettes should be filled with solution.

3. Prepare two solutions as indicated by the trial #1 row in the table below (do not mixthem!):

#1 - Into Erlenmeyer Flask #2 - Into Test Tube

Trial # A1 A3 A5 H2O A2 A4 Starch

1 0.0* 10 1 19 0 10 10

2 0 10 1 19 2 8 10

3 0 10 1 19 4 6 10

4 0 10 1 19 6 4 10

5 2 8 1 19 0 10 10

6 4 6 1 19 0 10 10

7 6 4 1 19 0 10 10

* amounts of added solution are to be taken as in units of mL

4. While the solutions for this trial are being prepared, the hot plate can be setup as depictedin Figure 3.2 and the stopwatch prepared for use.

5. Once a stable temperature has been established (do not turn the heating function on),record this temperature in your data sheet, immerse the Erlenmeyer flask into the waterbath and when ready, quickly add the contents of the test tube to the flask. Start timingabout halfway into adding the contents of the test tube.

6. You will notice that the solution is not coloured in any way, yet appears slightly cloudy. Stop timing when the solution is purple in colour (i.e., continue timing if it is only faintlypurple).

7. After the trial, clean and dry the test tube and flask as best as you can, then transfer theitems into the oven to eliminate any residual water.

8. Repeat steps as necessary until you have time measurements for all 7 trials.

34

C. The Effect of Temperature on ν01. Fill the 1 L beaker about 3/4 of the way with de-ionized water. This beaker will serve as

the new water bath for the remainder of this experiment.2. Heat the bath until the temperature is around 35 EC and then remove the beaker from the

heat source. Prepare the solutions using the same composition as outlined in trial #1 andhence, conduct trial #1 once again at this temperature. Record the initial temperature, theelapsed time and the final temperature (the temperature after colour change) on your datasheet. The average of the starting and final temperatures will be used as “thetemperature” at which the trial was conducted.

3. Repeat step C2, but this time heat the bath until the temperature is around 45 EC.

CalculationsPart B.1. For each trial in Part B, determine the initial concentration of all the ionic species present

and tabulate your results. You might find it useful to explicitly write down all of thechemical dissociation equations. Determine the average reaction temperature.

2. Determine the initial reaction rate, ν0, for all trials in part B. Express these rates in unitsof M min-1.

3. Create two plots, one of log ν0 versus log [I–] using the data from trials 1 – 4 and one oflog ν0 versus log [S2O8

2-] using the data from trials 1 and 5 – 7. Fit the data using afunction of the form y = mx + b. From the slope, determine the reaction order withrespect to each component. Round your value to the nearest half-integer.

4. Determine the rate constant for each trial and average these values. This mean value willbe known as the average rate constant at room temperature, krt.

test tube

Erlenmayer flask

water bath

stand

thermometer

hot plate / stirrer

stir bar

Figure 3.2

35

5. Determine the ionic strength of one of the solutions used and using equation [5],determine the rate constant at infinite dilution, k0, and compare this value with theliterature. Assume that the Debye-Hückel limiting law holds.

Part C.1. Determine the rate constant of the experiment at the two elevated experimental

temperatures. Use these two values and krt to create a plot of ln k versus 1/T anddetermine Ea for this reaction. Interpolate to find k298 (i.e., the rate constant at T = 298 K).

2. Using k298 and equations [7] and [8], determine ∆‡GE and ∆‡HE and comment upon thechemical significance of these values. Note: pay close attention to units when solving [7]for ∆‡GE. Use these two values to determine ∆‡SE and comment upon the physicalmeaning of the value isolated.

3. Using an appropriate reference (see list below), determine the standard Gibbs energy ofreaction, ∆rG°, from the ∆fG° values associated with each reacting species.

4. Using the values determined in steps C2 and C3, sketch a Gibbs energy reaction profile asthe reactants are converted into products.

Lab Questions 1. Starting from the form of the rate equation given in [2], show why a plot of log ν0 vs. log

[A] would give the order of reaction with respect to that particular reactant.2. Why were the NH4Cl and KCl salts used? If water were to be substituted for them, how

would it affect the observed rate constant and by how much? (find the percent difference) Hint: assume the same starting composition for one of the trials where a non-reagent saltwas used, but replace the added salt with water.

3. If reaction [3] is the third-order reaction, what are the units of k(T) [3]?4. An investigator is conducting kinetic experiments on a first-order reaction A P for

which the rate coefficient at 298 K is k0 = 0.001 s-1. The activation energy for the reactionis found to be 40 kJ mol-1. Assume that a simple Arrhenius expression gives thedependence of k upon temperature. The investigator runs the reaction in a containerwhose temperature is increased to 500 K. What is the change in the half-life time?

References1. Atkins, Peter and Julio de Paula. Physical Chemistry. 7th ed. New York: W. H.

Freeman, 2002. 256-258, 862-870, 879-881, 951-952, 956-963.2. Bernasconi, Claude F., ed. Investigation of Rates and Mechanisms of Reactions. 4th ed.

Toronto: John Wiley & Sons, 1986. 14-24.3. Wagman, Donald D. Selected Values of Chemical Thermodynamic Properties.

Washington: National Bureau of Standards, 1965.4. Howells, W. J. Journal of the Chemical Society (Resumed). 1939, 463-466.5. Amis, Edward S., and James E. Potts. J. Am Chem. Soc. 1941, 63, 2883-2888.6. Moews, P.C., and R. H. Petrucci. Journal of Chemical Education. 1964, 41, 549-551.

36

Data SheetA.Concentration of NH4Cl solution: ________ mol L-1

Concentration of KCl solution: ________ mol L-1

Concentration of (NH4)2S2O8 solution: ________ mol L-1

Concentration of KI solution: ________ mol L-1

Concentration of Na2S2O3 solution: ________ mol L-1

B.Trial # Temperature (EC) Elapsed Time (m:s)

1

2

3

4

5

6

7

Average Reaction Temperature: ________ EC

C.Trial Ti (EC) Tf (EC) T& (EC) Elapsed Time (m:s)

ca. 35 EC

ca. 45 EC

37

Sn 2%(aq) % 2Fe 3%

(aq) ! Sn 4%(aq) % 2Fe 2%

(aq) [1]

(A) 2Fe 3%(aq) ! 2Fe 2%

(aq) % 2e –

(B) Sn 4%(aq) ! Sn 2%

(aq) % 2e –[2]

EXPERIMENT 4THERMODYNAMIC FUNCTIONS of a GALVANIC CELL

IntroductionChemical Reactions Involving the Transfer of Electrons

Numerous chemical reactions have been studied which involve the transfer of electronsfrom one species to another. The chemical species which has donated electrons is defined as thereducing agent and is said to have undergone the process of oxidation over the course of thereaction. The chemical species which receives the electrons is defined as the oxidizing agent andis said to have undergone the process of reduction over the course of the reaction. A redox(reduction/oxidization) reaction is the term applied to a reaction in which electrons aretransferred from one chemical species to another.

Half-ReactionsAlthough the processes of oxidation and reduction occur simultaneously, the technique

most widely used for understanding redox reactions involves separating them into half-reactions. This separation can be done in more than one equivalent fashion, with one way being outlinedbelow. Consider the redox reaction where tin(II) and iron(III) are the reagents:

Above is the net ionic equation describing this redox reaction in aqueous media; below, the sameredox reaction is expressed as two reduction half-reactions:

By taking the difference between the two half-reactions in [2], one can arrive at [1].

The Electrochemical CellThe separation of redox reactions into half-reactions is convenient, and it reminds one of

the electrochemical cell, a common physical construct used when studying this type of reaction. Each physical half-cell is composed of a single electrode, which is placed in contact with anelectrolyte of some sort. A typical half-cell will have a strip of metal, M, (the electrode) which isin contact with an aqueous solution that contains ions of this metal (the electrolyte), Mn+, where nis the charge number associated with a particular oxidation state of the metal. A low resistancewire will allow the passage of electrons from one electrode to the other. Note that there are manyvariations to this scheme, one of which will be utilized in this experiment. The fashion in whichthe cell circuit is completed also varies, but one regularly used method involves a salt bridge, andit is the method that will be used for this experiment. A salt bridge is often made by adding agarto an electrolytic solution followed by heating. Upon cooling, the solution will turn into a gel. Aconnected set of two half-cells is called an electrochemical cell; a general schematic of which isprovided below:

38

µ B & µ A ' ∆r G ' we,max ' – νöEcell [3]

Once connected, a redox reaction may begin to occur spontaneously, accompanied withthe production of electricity. In cases of this sort, the electrochemical cell is referred to as agalvanic cell. If electricity is required to drive the chemical reaction, the redox reaction beingconsidered is non-spontaneous and the electrochemical cell is known as an electrolytic cell. Regardless of the type of electrochemical cell, the flow of electrons is from the anode (this is theelectrode at which the process of oxidation occurs) to the cathode (the electrode at whichreduction occurs).

Gibbs Energy Change and The Cell PotentialIn galvanic cells, as with any spontaneous process, the redox reaction occurs

spontaneously due to the fact that the chemical potential, µ, of the system decreases. If thesystem is under constant temperature and pressure conditions, the chemical potential change ofthe system can be related in a straightforward fashion to the Gibbs energy change, ∆rG. Additionally, as long as the process is reversible, ∆rG can be related to the cell potential, Ecell, (orthe electromotive force, (emf)) which is defined as being the electrical potential differencebetween the anode and cathode and is related to the maximum amount of electrical work, we,max,that can be done by the system. The above statements are represented algebraically:

Where B is some final state, A is some initial state, Ecell is in units of V mol-1, ν is the number ofelectrons transferred (units of mol), and ö is the Faraday constant (96 485 C mol-1).

Figure 4.1

39

∆ r G ' ∆r GE % RT ln Q [4]

Ecell ' EcellE &RTνö

lnQ [5]

dEcell

dT'

∆ r Sνö

[6]

Ecell as a function of Cell CompositionThe relationship between the reaction quotient, Q, and the Gibbs energy of reaction takes

on the form below:

Where ∆rGE is the change in Gibbs energy when the system is under standard conditions. Theabove equation can be re-expressed using [3]:

The form above is known as the Nernst equation and is defined as the standard emf.EcellE

Other Thermodynamic FunctionsIt has been established that the emf is related directly to the Gibbs energy of reaction, but

the fashion in which the emf varies with respect to temperature can allow one to determine thechange in another thermodynamic function of state. By taking the derivative of the isolated Ecellvalue with respect to temperature, the entropy of reaction, ∆rS, can be determined:

Once the values for ∆rS and ∆rG are known, the enthalpy of reaction, ∆rH, for theparticular redox reaction may be determined.

Materials

• 2 – 50 mL volumetric flasks• 2 – scoopula• 1 – 50 mL graduated cylinder• 1 – hot plate/stirrer combo

w/ temperature probe• 1 – glass stir rod• plastic tubing• 4 – 100 mL beakers• 2 – 100 mL beaker lids• 1 – magnetic stir bar• 1 – Pasteur pipette• 1 – water bath

• 1 – 10.0 mL pipette• 1 – 1 MΩ resistor• 1 – voltmeter• 1 – platinum electrode• 1 – zinc electrode• potassium ferrocyanide

(K4Fe(CN)6C3H2O)• potassium ferricyanide (K3Fe(CN)6)• zinc chloride (ZnCl2)• pH 4 buffer• potassium nitrate (KNO3)• agar

ProcedureA. Preparation of the Salt Bridge (another group member should start on part B)1. Using the 50 mL graduated cylinder, measure out 50 mL of de-ionized water and transfer

this to a clean 100 mL beaker.2. To this beaker, first add enough KNO3 to make a 0.1 M solution, then add 2 grams of

40

Vtemperature probe

hot plate/stirrer

100 mL beaker w/ Fe(CN)63-/4-

100 mL beaker w/ Zn2+

Pt electrode connectionZn electrode connection

34.5

resistor/voltmeter

Figure 4.2

agar.3. Heat this solution with stirring to around 85 EC. To do this, turn the heating function of

the hot plate on, and set the dial to around 150 EC. On the temperature probe display, setthe temperature to 95 EC (note that this is about 10EC above the target temperature). The probe is designed such that the solution may be heated by the hot plate to a range ofdesired temperatures. Wait until the solution is clear and bubbles have started to form.

4. Remove the beaker from the heat, then using a Pasteur pipette attached to one end of thetubing, withdraw a portion of the hot solution into the tubing. Once this cools (let sitabout 30 – 40 min), it will form a gel and may be used as the salt bridge.

B. Preparation of Required SolutionsPrepare the following solutions using de-ionized water (note: do not use tap water):1. 50 mL of a solution that is 0.10 M in BOTH potassium ferrocyanide (K4Fe(CN)6C3H2O)

and potassium ferricyanide (K3Fe(CN)6).2. 50 mL of ca. 0.10 – 0.20 M zinc chloride (ZnCl2). Make sure to add 10.0 mL of pH 4

buffer before topping up with the de-ionized water.

C. Effect of Temperature on Ecell1. Create an ice slurry using the water bath as a container and then assemble the

electrochemical cell as shown in Figure 4.2. Add a stir bar to the bath as well.

2. Cut the ends of the salt bridge so that the gel runs the full length of the tubing. Whenready, turn the heating/stirring functions on and set the temperature dial to around175 EC. On the temperature probe display, set the temperature to 20 EC.

3. As the temperature gets close to the desired value (- 1 EC away), add the salt bridge.

41

4. Proceed to take a measurement of the cell potential at 10 EC and record this value in yourdata sheet. Note: the voltmeter should be set to read direct current, .&&&V

5. Once the reading has been made, remove the salt bridge and heat the system to 15 EC.6. Continue taking measurements and heating the system until the data sheet is completed.

Calculations1. Write out the chemical equation for the redox reaction that was under study in this lab

and use it to calculate Q. Note: you do not have to worry about activities. Then, using Qand [5], find the value of at each temperature point.EcellE

2. Create a plot of versus temperature.E Ecell

3. Using the line of best fit, interpolate what the value of should be at T = 298 K. ThisE Ecell

will be known as . Compare your value with the literature.E E298

4. Using the value isolated for , calculate ∆rGE.E E298

5. Using the plot composed in step C2, determine ∆rSE for this reaction. Then, using thevalue of ∆rGE calculated earlier, find ∆rHE.

Lab Questions 1. The pH 4 buffer was used to maintain the cell potential over the course of the reaction.

What possible side reaction does the use of this buffered solution prevent and why wouldthis cause a change in the cell potential? (see references)

2. Distinguish cell potential and electromotive force.3. Using the value ∆fGo(Al(aq)

+3) = -481.2 kJ/mole, calculate ∆fGo in kJ mol-1 for Ba(aq)+2

using the following Galvanic cell: Al|Al(NO4)3||Ba(NO3)2|Ba4. Based on the electrochemical process:

3Pb(aq)+2 + 2Al(s) 2Pb(s) + 2Al(aq)

+3 Calculate the cell voltage at 25 oC given starting concentrations: [Al(aq)

+3]0 = 0.00300 Mand [Pb(aq)

+2]0 = 2.45 M.5. A student obtained a negative reading in the cell potential during the experiment. Based

on that, the student concluded that the cell reaction is not spontaneous. Is the student'sjudgment correct? Why?

References1. Atkins, Peter and Julio de Paula. Physical Chemistry. 7th ed. New York: W. H.

Freeman, 2002. 262-281.2. Brown, Theodore L., LeMay, H. E., and Bruce E. Bursten. Chemistry: The Central

Science. 6th ed. Englewood Cliffs: Prentice Hall, 1994. 720-743.3. McSwiney, Harry D. Journal of Chemical Education. 1982, 59, 165.4. Probst, Donald A., and Giles Henderson. Journal of Chemical Education. 1996, 73, 962-

964.

42

Data SheetB.Actual concentration of ZnCl2 solution: ________ mol L-1

C.Temperature (EC) Ecell (V mol-1)

10

15

20

25

30

35

40

45

50

43

N ' Kc a (1)

EXPERIMENT 5ADSORPTION FROM SOLUTION

IntroductionThe term adsorption is used to describe the fact that there is a greater concentration of the

adsorbed molecules at the surface of the solid than in the bulk solution. In general, one usessolid adsorbents of small size and often with surface imperfections such as cracks and holeswhich serve to greatly increase the surface area per unit mass over the apparent geometrical area.Such small, porous particles may have specific areas in the range from 10 to 1000 m2 g-1. Someexamples of adsorbents commonly used in experiments of this kind are charcoal, silica gel,alumina, zeolites, and molecular sieves. The adsorption from aqueous solutions of acetic acid oncharcoal will be investigated in the present experiment.

The type of interaction between the adsorbed molecule and the solid surface varies over awide range, from weak nonpolar van der Waals' forces to strong chemical bonding. Examples ofadsorption where ionic or covalent bonding occurs include the adsorption of chloride ions andsilver chloride (ionic) or of oxygen gas on metals where oxygen-metal bonds are formed(covalent). In these cases, the process is called chemisorption (chemical adsorption) which isgenerally characterized by high heats of adsorption (from 10 to 100 kcal mol-1 of gas adsorbed). Chemisorption is highly specific in nature and depends on the chemical properties of both thesurface molecules and the adsorbed molecules. Adsorption arising from weaker van der Waals'and dipole forces is not specific in character and can take place in any system at low or moderatetemperatures. This type of adsorption is called physisorption (physical adsorption) and is usuallyassociated with low heats of adsorption (less than 10 kcal mol-1). Physical adsorption forces aresimilar to those which cause condensation of gases into liquid or solids. When an adsorbingmolecule approaches the surface of the solid, there is an interaction between that molecule andthe molecule in the surface which tends to concentrate the adsorbing molecules on the surface inmuch the same way that a gas molecule is condensed onto the surface of bulk liquid. Anotherrespect in which physical adsorption is similar to liquid condensation is the fact that molar heatsare of adsorption are of the same order of magnitude as molar heats of vaporization.

The amount adsorbed per gram of solid depends on the specific area of the solid, theequilibrium solute concentration in the solution (or pressure in the case of adsorption from thegas phase), the temperature, and the nature of the molecules involved. From measurements atconstant temperature, one can obtain a plot of N, the number of moles adsorbed per gram ofsolid, versus c, the equilibrium solute concentration. This is called an adsorption isotherm.

Often, it is possible to represent experimental results over a limited range by an empiricalisotherm suggested by Freundlich:

where K and a are constants which have no physical significance, but can be evaluated by a plotof log N versus log c. However, eq. (1) fails to predict the behaviour observed at low and highconcentrations. At low concentrations, N is often directly proportional to c, whereas at highconcentrations N usually approaches a constant limiting value which is independent of c.

44

θ 'kc

1 % kc (2)

cN

'c

Nm

%1

kNm(3)

A ' NmN0σ × 10&20 (4)

Much effort has been devoted to developing a theory of adsorption which would explainthe observed experimental facts. In some simple systems, a theory derived by Langmuir can beapplied. This theory is restricted to cases where only one layer of molecules can be adsorbed atthe surface. Monolayer adsorption is usually observed in the case of chemisorption from the gasphase or adsorption from solution, and is distinguished by the fact that the amount adsorbedreaches a maximum value at moderate concentrations (corresponding to complete coverage of thesurface of the adsorbent by a layer one molecule thick) and remains constant with furtherincreases in concentration. The Langmuir isotherm can be derived from either kinetic orequilibrium arguments and is most commonly applied to the chemisorption of gases. We shallgive a form appropriate to adsorption from solution:

where θ is the fraction of the solid surface covered by adsorbed molecules and k is a constant atconstant temperature. Now θ = N/Nm, where N is the number of moles adsorbed per gram ofsolid at an equilibrium solute concentration, c, and Nm is the number of moles per gram requiredto form a monolayer. Making this substitution and rearranging Eq. (2), we obtain

If the Langmuir isotherm is an adequate description process, then a plot of c/N versus c will yielda straight line with slope 1/Nm. If the surface area, σ, occupied by an adsorbed molecule on thesurface is known, the specific area, A, (in square meters per gram) is given by

where N0 is Avogadro's number and σ is given in square angstroms.

Materials6 250 ml Erlenmeyer flasks with glassstoppers or rubber stoppers2 100 ml Erlenmeyer flasks (for titration)6 100 ml beakers (for filtering)1 250 ml beaker6 100 ml volumetric flasks1 Funnel holder (or three rings, with clampsand stands)3 funnels

Fine porosity filter paper50, 20, and 10 mL pipettes2 50 ml burettesThermometer2 bulbsSelf adhesive labelSpatula

The following stock solutions will be provided: 1. 0.30 M acetic acid2. 0.20 M sodium hydroxide3. Phenolphthalein indicator4. Activated charcoal

45

Procedure1. Clean and dry six 250 ml Erlenmeyer flasks. These should either have glass stoppered

tops or be fitted with rubber stoppers. Label the flasks from 1 to 6. Place approximately1 g of charcoal in each flask (the weight need not be exactly 1 g but it should be known tothe nearest milligram).

2. To each flask, add 100 ml of acetic acid solution, measured accurately with a pipette.Suggested initial concentrations are 0.30, 0.27, 0.24, 0.21, 0.18 and 0.15 M. Using 100mL volumetric flasks, prepare the following solutions by adding the respective quantitiesof 0.30 M acetic acid stock solutions and diluting up to the mark using de-ionised water(in the case of flask 1, fill with acetic acid to the mark):

Flask Activated carbon(g) Acetic Acid conc. Acetic acid (mL) De-ionized H2O (mL)1 0.30 100 02 0.27 90 103 0.24 80 204 0.21 70 305 0.18 60 406 0.15 50 50

Add these solutions to the six flasks containing the charcoal adsorbent.3. Close the flasks with stoppers, and shake them periodically for a period of 30 min. After

you are finished shaking, allow at least 20 minutes to reach equilibrium.4. While waiting for the solution to equilibrate, determine the exact concentration of the

stock acetic acid solution by titrating two 10 ml samples of the stock solution with 0.20M NaOH using phenolphthalein as indicator.

5. After equilibrium has been reached, measure the temperature of solutions. Filter all thesamples through fine filter paper. Discard the first 10 mL of the filtrate as a precautionagainst adsorption of the acid by the filter paper.

6. After filtering, titrate samples of the filtrates with 0.20 M NaOH to determine theequilibrium concentration of acetic acid. You will find it convenient to titrate accordingto the following scheme:

Filtrate No. Sample Volume (mL) Number of Samples Volume of 0.20 NaOH (mL)1 10 22 10 23 25 24 25 25 30 26 30 2

Titrate all solutions accurately!

46

CalculationsCalculate the final concentration of acetic acid for each sample. From the values of the

initial (c0) and final concentrations (ce) of acetic acid in 100 mL of solution, calculate the numberof moles present before and after adsorption, and obtain the number of moles adsorbed bydifference. Compute N, the number of moles of acid adsorbed per gram of charcoal. Plot anisotherm of N versus the equilibrium (final) concentration c in moles per litre.

As suggested by Eq. (3), plot c/N versus c. Draw the best straight line through thesepoints, and calculate Nm from the slope. On the assumption that the adsorption area of acetic acidis 21 Å2, calculate the area per gram of charcoal from Eq. (4).

Lab Questions1. What are three assumptions in Langmuir isotherm?2. What is a BET isotherm?3. M. G. Olivier and R. Jadot (J. Chem. Eng. Data 42, 230 (1997)) studied the adsorption of

butane on silica gel. They report the following amounts of adsorption (in moles perkilogram of silica gel) at 303 K:

p/kPa 31.00 38.22 53.03 76.38 101.97 130.47 165.06 182.41 205.75 219.91

n/(mol kg-1) 1.00 1.17 1.54 2.04 2.49 2.90 3.22 3.30 3.35 3.36

Fit these data to a Langmuir isotherm, and determine the value of Nm that corresponds tocomplete coverage and the constant k.

References1. Shoemaker, David P., Garland, Carl W., Steinfeld, Jeffrey I., and Niebler, Joseph W.

Experiments in Physical Chemistry, 4th ed. New York: McGraw-Hill, 1981. 332-337.2. Atkins, Peter and Julio de Paula. Physical Chemistry, 7th ed. New York: W. H. Freeman,

2002. 987-994.