Chemistry 59-240

CHEMISTRY 59-240

PHYSICAL CHEMISTRY LABORATORY MANUAL

FALL 2010 10th Edition - Version 1.3 DEPARTMENT OF CHEMISTRY & BIOCHEMISTRY UNIVERSITY OF WINDSOR

TABLE OF CONTENTS

GENERAL INSTRUCTIONS

Emergency Procedures...... 3 Safety Regulations & Quiz...... 9 Policy on Plagiarism...... 21 Student Contract...... 25 Marking Scheme and Outline...... 27 Medical Certificate ...... 31

EXPERIMENTS

Experiment 1: Determination of ΔcH: Bomb Calorimetry...... 33 Experiment 2: Vapour Pressure of Pure Liquids...... 41 Experiment 3: Surface Tension of n-Butanol and Amount Adsorbed...... 49 Experiment 4: Heat of Reaction in Solution: Constant Pressure Calorimeter...... 55 Experiment 5: Liquid-Vapour Equilibrium in a Binary System...... 59

Emergency Procedures

EMERGENCY PROCEDURES

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 Wash Stations located in each hallway, 173-3 (Lab E), 173-6 (Lab F) Safety Showers located in each laboratory

EMERGENCY PROCEDURES

Discovery of a Fire

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

5. Remain calm and to stay away from the fire.

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

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

Emergency Procedures

Sounding of Evacuation (Fire) 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 hot plates.

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 roles in building evacuation and direct you the Assembly Area.

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

Critical Reminders:

9. Do Not Use Elevators. Only authorized persons including the City of Windsor Fire Department, Campus Police, and Fire Safety personnel shall be permitted to use the elevators during fire emergencies.

10. Firefighting is the responsibility of the City of Windsor Fire Department. If the fire is small or in its earliest stages and can be fought effectively with available fire extinguishers, then you may attempt to extinguish such fires providing there is no life safety hazard to the user and such action will not endanger others.

Your student safety is the most critical factor, not extinguishing the fire itself!

Medical Emergency

Dial 911 to activate emergency services.

If there is an emergency in which someone’s life is at stake and the loss of Important: a minute or two is critical, then it is appropriate to pull the fire alarm.

Emergency Procedures

FIRE SAFETY & FIRE EXTINGUISHERS

Introduction

While proper procedure and training can minimize the chances of an accidental fire, you must still be prepared to deal with a fire emergency should it occur. This document teaches you the basics about fire extinguishers – proper types, how to use them, when and when not to use them as well as the proper procedures to follow should a fire occur. It is not a comprehensive guide, be sure to read the DISCLAIMER given below.

If yours or your student’s clothing is on fire (and the floor is not), STOP, DROP, and ROLL on the ground to extinguish the flames. If you are within a few feet of a safety shower or fire blanket, you can use these instead, but do not try to make it “just down the hall” if you are on fire. If one of your coworkers or students catches fire and runs down the hallway in panic, tackle them and extinguish their clothing.

Types of Fires

Type of Fire Materials Involved

Ordinary materials like burning paper, lumber, cardboard, plastics, Class A etc.

Flammable or combustible liquids such as gasoline, kerosene, Class B and common organic solvents used in the laboratory

Energized electrical equipment, such as appliances, switches, Class C panel boxes, power tools, and hot plates and stirrers. Do not use water to extinguish due to the risk of electrical shock.

Combustible metals, such as magnesium, titanium, potassium, and Class D sodium as well as pyrophoric organometallic reagents such as alkyllithiums, Grignards, and diethylzinc. Do not use water or CO2 to extinguish due to potentially violent reactions. The above materials will burn at high temperatures and will react violently with water, air, and/or chemicals. Handle with care!

Some fires may be a combination of these. The fire extinguishers have an ABC rating and the higher the number, the more firefighting power it has. The classification of extinguisher in the laboratories is labelled “20 BC”.

Emergency Procedures

Types of Fire Extinguishers

Type of Extinguisher Materials Involved

Useful for class A (paper, etc.) fires only. Use on class Water B, C, or D fires will cause flames to spread or the hazard made greater.

Useful for class ABC fires. This extinguisher leaves a blanket of non-flammable material on the extinguished material which reduces the likelihood of reignition. There Dry Chemical are two types of dry chemical extinguishers: Type BC: Contains sodium or potassium bicarbonate Type ABC: Contains ammonium phosphate Useful for class B and C fires only. This extinguisher Carbon dioxide (CO2) leaves behind no harmful residues. Do not use on class A fires because the material usual reignites. This extinguisher is not approved for class D flammable metal fires such as Grignard reagents, alkyllithiums, and sodium metal because CO2 reacts with these materials.

Typical small laboratory fires can be easily controlled by the use of sand or with a dry chemical (ABC) extinguisher provided that you are properly trained to use one.

Using Fire Extinguishers

You are not required to fight a fire... Ever! If you have the slightest doubt about your control of the situation, DO NOT FIGHT THE FIRE. Please see the DISCLAIMER below.

1. Use a mental checklist to make a Fight-or-Flight Decision. Attempt to use an extinguisher only if ALL of the following apply:

. The building is being evacuated (fire alarm is pulled) . The fire department is being called (dial 4444 to activate 911 on campus) . The fire is small, contained and not spreading beyond its starting point. . The exit is clear, there is no imminent peril, and you can fight the fire with your back to the exit. . You can stay low and avoid smoke. . The proper extinguisher is immediately at hand. . You have read the instructions and know how to use the extinguisher.

IF ANY OF THESE CONDITIONS HAVE NOT BEEN MET, DON’T FIGHT THE FIRE YOURSELF. CALL FOR HELP, PULL THE FIRE ALARM AND LEAVE THE AREA.

Emergency Procedures

2. Whenever possible, use the “Buddy System” to have someone back you up when using a fire extinguisher. If you have any doubt about your personal safety, or if you cannot extinguish a fire, leave the building but contact a firefighter to relay whatever information you have about the fire.

3. Operating the fire extinguisher:

Stay low behind the fire extinguisher as the fire usually flashes back over the top of the extinguisher upon the first burst of air. Then,

P pull pin A aim low at base of flame S squeeze the lever S sweep from side to side

. Do not walk on area that you have “extinguished” in case the fire reignites or the extinguisher runs out! Remember, you usually can’t expect more than 10 full seconds of extinguishing power on a typical unit and this can be significantly less if the extinguisher was not properly maintained or partially discharged.

. Again, proper training is required by provincial or federal laws.

4. Recharge any discharged extinguisher immediately after use. If you discharge an extinguisher (even just a tiny bit) or pull the pin for any reason, inform the Laboratory Coordinator to arrange a replacement.

DISCLAIMER: These pages contain guidelines for the use of fire extinguishers and are not meant to be a comprehensive reference. There are many circumstances that these guidelines can not foresee and you should recognize the inherent danger in relying solely on this information!

Safety Rules

SAFETY RULES IN THE LABORATORY

Safety is a matter of the greatest importance to everyone in the Chemistry Department. The principal danger is fire, but others (toxic fumes, sharp objects, etc.) are also of concern. The following sets out regulations, which apply to all students in the teaching laboratories. Strict adherence to these is a MUST. Defaulters may be excluded from the laboratory and course.

The following rules are strictly enforced:

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

. Safety glasses or goggles must be worn at all times in the laboratory, including when washing glassware. Once an experiment begins, glasses must be donned and must remain on until the student exits the laboratory. Students refusing to wear or are not in possession of safety glasses will be refused permission to perform the experiment. Contact lenses must not be worn in the laboratory.

. It is mandatory that a laboratory coat be worn at all times during the experiment.

Safety goggles and laboratory coats are available for rental. The cost of the rental(s) will be deducted from the student’s caution card. Rental fees: $5 – laboratory coats $2 – safety goggles

. Laboratory work is only permitted on assigned laboratory periods in the assigned laboratory room and in the presence of a GA. Unauthorized experimental work conducted outside of stipulated laboratory hours is prohibited and forbidden by departmental regulations on grounds of safety. Severe disciplinary action will be taken against anyone attempting unauthorized experiments in the laboratory.

. Students must ask for permission to leave the laboratory premises. Failure to notify the GA will result in a 10% grade deduction on the laboratory report for the experiment being performed.

. Visitors are not permitted to enter the laboratory during scheduled laboratory periods.

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

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

. Open-toed shoes or sandals may not be worn. Students not wearing the proper footwear will neither be permitted to perform the scheduled experiment.

. Long hair must be tied back at all times.

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

Safety Rules

. Transfer harmful reagents in the fume hood to eliminate the dispersion of toxic fumes throughout the laboratory. Always wear gloves as protection.

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

. Never place your face directly over the top of a container which is being heated or after the heat is removed. Contents may “bump” and be violently ejected from the container.

. Never dispose of glass or sharp objects – such as Pasteur pipets – into the regular garbage containers. Always use the specifically designated and labelled containers for (broken) glass or sharp objects located in the back of the laboratory.

. Never dispose of ANY chemicals into the sink and down the drain. Use ONLY the properly labelled waste containers located in the waste fume hood for disposal of liquid and solid chemicals. Check the label of all waste containers twice before adding chemical waste to the container(s).

. Keep working space clean and free of apparatus and/or other materials, especially when flammable materials are in use. Wipe up spilled materials immediately.

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

All injuries must be reported, no matter how serious, to the GA(s). The GA Important: will provide instruction on further first aid, and notify the Laboratory Coordinator.

Conduct, House Rules, and Safety

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

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

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

Chemical Safety

CHEMICAL SAFETY IN THE LABORATORY

Chemical Reagents

All of the reagents needed to conduct an experiment are provided in the fume hood(s).

1. Solid Reagents Balances for general laboratory work are located within the laboratory on the bench at the front of the laboratory. Most solid reagents will be located next to the balances. Any spill chemicals on or around the balance will be the student’s responsibility to clean up.

2. Hazardous Reagents - Acids and Bases Hazardous chemicals, such as concentrated acids and bases, solvents, noxious chemicals, and other hazardous or volatile substances, are found in the fume hood(s). Clean up any acid or base spills to prevent a possible chemical burn caused by these reagents.

3. Liquid Reagents

Any liquid reagents, if noxious or smelly, will also be found in the fume hood(s). Always replace caps on containers immediately after use. An open container is an invitation for a spill. Furthermore, some reagents are very sensitive to moisture, and may decompose if left open.

4. Avoid Contamination of Chemicals Do not put chemicals back into reagent bottles; returning any unused chemical to its original container risks contamination. Extra reagents must be placed in the appropriate chemical waste container. If possible, take only the amount that is needed to prevent waste.

5. Personal Safety with Chemicals  Avoid direct contact with any chemical. Always wear gloves when handling chemicals.  Keep laboratory chemicals off hands, face, and clothing (including your shoes).  Never smell, inhale, or taste laboratory chemicals. Be sure there is adequate ventilation.  Wash hands thoroughly with soap and water after handling any chemicals, especially before leaving the laboratory.

Chemical Safety

Chemical Hazards

The following chemical hazard terms listed below are words that one may stumble upon when viewing a chemical MSDS. These are included to better understand MSDS sheets, not because they will apply to specific compounds used in laboratory.

1. Flammable compounds have a flash point below 100 °F (37.8 °C), and hence, may ignite and burn.

2. Corrosive or caustic compounds cause obvious damage to living tissue. Corrosives act either directly, by chemically destroying the part (oxidation), or indirectly by causing inflammation. These chemicals will cause damage (immediate burns) to your skin:  Corrosive: usually applies to acids  Caustic: usually applies to bases

3. Strong oxidizer reactions are usually very exothermic (give off heat). Therefore, oxidizers can cause other materials to combust more readily (or upon contact!) or make fires burn more fiercely. Oxidizers are extremely reactive.

4. Volatile compounds have a high vapour pressure and easily form vapours at normal temperature and pressure. The vapour could be flammable or toxic, or both.

Chemical Toxicity

MSDS may include several terms related to chemical toxicity, which is related to the adverse effects of a chemical on a living system. There are two types of toxicity:

Acute toxicity: The chemical may have a rapid bodily absorption and can exert an effect during a single exposure

Chronic toxicity: The chemical may exert an effect because of repeated exposure over a period of time (days, months, years) and the exposures may be cumulative

MSDS (Material Safety Data Sheets) for all chemicals used in the Important: experiments can be found in a grey binder at the front of the laboratory by the safety shower.

The NFPA Diamond

The National Fire Protection Association (NFPA) has developed a hazard warning symbol called the NFPA Diamond, which provides quick information regarding a chemicals hazard profile.

The type of hazard (flammability, health, reactivity) is indicated by colour and location. Hazards are rated from 0 (very low hazard) to 4 (extreme hazard) in each category, with a provision for indicating any special hazards.

Chemical Safety

Hazard Levels

0 = Very Low 1 = Slight 2 = Moderate 3 = Severe 4 = Extreme

MSDS often abbreviate the NFPA Diamond Ratings as: H-F-R, i.e.; 4-3-2

The NFPA Diamond is often placed on reagent bottles in this form by many of the manufacturers. The sequence 4-3-2 refers to the Health (blue), Flammability (red), and Reactivity (yellow) ratings, respectively and is always in this order.

Specific Hazards

The white area of the triangle is reserved for specific hazards. Some of the special symbols used are shown below. Below, the "corrosives" hazard is shown in the NFPA diamond.

Chemical spills

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

Chemical spills on bench tops, fume hoods, and balance area should be cleaned up immediately. Failure to clean up a spill can result in a fellow student be unknowingly burned or exposed to toxic chemicals. Treatment with an appropriate neutralising reagent, if necessary, should be based on consultation with the GA and/or Laboratory Coordinator.

Chemical Safety

Fires

In the event of a fire in the laboratory, turn off all gas, and lower the sashes of the fume hoods if possible before leaving the room. Do not hesitate to shout “FIRE” or sound the building alarm for any sizable 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. Do not run about, which includes running any distance for a safety shower. Use the safety shower or fire blanket if very close by.

First Aid

The section below lists some first aid procedures for minor injuries. More serious injuries must be reported to the GA, as well as the Laboratory Coordinator.

1. Acid or Base Burns Rinse the affected area with large quantities of water for at least 15 minutes. Bases have a slippery feeling (like soap); acids cause a "non-skid" feeling and may burn. Rinse until the skin returns to normal. Do not be concerned with neutralization, just wash. If a sizable quantity of corrosive material is on the clothes, prevent skin contact by quickly removing the clothing. An eye wash is located directly outside of the laboratory for washing out injured eyes. Be sure to ask for help, no matter how slight the eye injury is.

2. Minor Cuts Internal contact with chemicals can occur through a cut from broken glass. Wash the wound well with water. If necessary, apply pressure to stop the flow of blood. The Laboratory Coordinator, who is trained in First Aid, can treat cuts. If shattered glass was involved, inspect for fragments still in the wound before bandaging.

3. Minor Burns Immerse the burned area in cold water until the pain is alleviated. Use of salves or ointments is discouraged.

4. Accidental Ingestion Report at once if any chemicals may have entered the mouth, whether or not if any chemical has been swallowed. Delay will not make any necessary treatment less unpleasant or less necessary. Do not drink anything unless instructed by a medical professional.

Safety Quiz

SAFETY QUIZ

Please take a few moments to answer these questions on the Answer Key provided.

Question and Answer: 1. Where is the location of the fire alarm pull? 2. How many fire extinguishers are in the laboratory? Where is its/their location(s)? 3. Where is the nearest eyewash? 4. Where is the safety shower in the laboratory? 5. Where is the fire blanket in the laboratory?

Multiple Choice:

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

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

8. Why should contact lenses never be worn in the laboratory? (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

9. Why should you wear safety eyewear as soon as you walk into the laboratory and at all times while in the laboratory even if you personally are not working with any chemicals? (a) Discomfort is a part of science (b) Someone at another workstation may be working with chemicals and splash you (c) They make you look smart

10. Which of the following statements is not true upon entering the laboratory: (a) Laboratory work can be started immediately even if the GA is not yet present (b) Safety eyewear and lab coat must be worn prior to moving to the work bench/station (c) All electronic devices must be turned off and left by the area where bags and coats are designated (not in lab coat pockets)

11. The main routes of exposure to chemicals are: (a) Through open-toed shoes (b) Inhalation, ingestion and skin/eye contact (c) Eyes, nose, mouth and ears 15

Safety Quiz

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

Where is this information located? (a) At the sink (b) In the GA locker (c) In a wall container at the front of the room

13. Why should broken or chipped glassware never be used? (a) It can explode when heated or under pressure (b) Glass cuts are extra dangerous in laboratory due to chemicals (c) Broken or chipped glassware can always be used, as long as it doesn’t leak

14. 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 laboratory (c) In the back of someone else’s locker when they are not looking

15. An excess of a chemical that has been measured out should be: (a) disposed of as waste (b) returned to the bottle (c) flushed down the drain

16. 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

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

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

19. 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 GA (b) Call 911 (c) Run madly about the laboratory until it dries

20. 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) Nothing, unless the chemical causes discomfort (c) Rinse under the safety shower for 5 minutes

Safety Quiz

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

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

Identify the WHMIS symbols by choosing the letter of the symbol which corresponds with the descriptions of the classification.

23. Biohazardous (a)

24. Toxic (b)

25. Poisonous (c)

26. Flammable (d)

27. Oxidizing (e)

28. Corrosive (f)

29. Dangerously reactive (g)

30. Compressed gas (h)

Safety Quiz

Student Name ______

Student No. ______Laboratory Section ______

Question and Answer:

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

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

3. Where is the nearest eyewash?

4. Where is the safety shower in the laboratory?

5. Where is the fire blanket in the laboratory?

Multiple Choice:

6. 11. 16. 21. 26.

7. 12. 17. 22. 27.

8. 13. 18. 23. 28.

9. 14. 19. 24. 29.

10. 15. 20. 25. 30.

Submit this copy to your GA.

Policy on Plagiarism

POLICY ON PLAGIARISM

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

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

ACADEMIC INFORMATION Undergraduate Degree Regulations

2.4.22 POLICY ON PLAGIARISM

Plagiarism is defined as: "The act of appropriating the literary composition of another, or parts of passages of his or her writing, or the ideas or language of the same, and passing them 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 and all 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 work in preparing essays, papers, reports, theses and publications. It is imperative that both the data and ideas obtained from any and all published or unpublished material be properly acknowledged and their sources disclosed. Failure to follow this practice constitutes plagiarism and is considered to be a serious offence. Thus, anyone who knowingly or recklessly uses the work of another person and creates an impression that it is his or her own, is guilty of plagiarism.

Plagiarism also includes submitting one's own essay, paper, or thesis on more than one occasion. Accordingly, it is expected that a thesis, essay, paper or a report has not been and is not concurrently being submitted for credit for any other course. In exceptional circumstances and with the prior agreement of the instructor, a student may use research completed 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 website dedicated to student integrity at http://www.uwindsor.ca/aio

Policy on Plagiarism

SENATE BY-LAW 31

UNIVERSITY POLICY IN RESPECT TO JUDICIAL PROCEDURES ARTICLE I. SANCTIONS AND DEFINITIONS

Proscriptions Stated University discipline is limited to student misconduct which adversely affects the University community's pursuit of its educational objectives. Students are expected to conduct themselves in a manner compatible with the objectives and purposes of the University of Windsor. Any student at the University of Windsor whose conduct is improper in that it exhibits a lack of integrity touching upon the educational objectives and requirements of the University must be disciplined appropriately in the interest of safeguarding and upholding these standards.

It is desirable to define and identify further the standards demanded of each student at the University of Windsor in the interest of educational integrity. Enumerated below are illustrations of improper conduct which would lead to an inference of lack of integrity. These are illustrative only and shall not be taken as in any way limiting the generality of the high standards 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 defined as follows: a. Dishonesty, such as cheating, plagiarism, impersonation at an examination, or knowingly furnishing false information to the University. b. Forgery, alteration, or use of University documents, records, or instruments of identification with intent to defraud. c. Intentional obstruction or disruption of teaching, research, administration, disciplinary proceedings, or other University activities, including public service functions, and other authorized activities on University premises. d. Malicious abuse of any person on University premises or at University sponsored or University supervised functions or malicious conduct which threatens, endangers or harasses any such person. e. Theft from or deliberate damage to University premises or theft of or deliberate damage to property of a member of the University community on University premises. f. Failure to comply with directions of members of the University administration or of the teaching staff acting in the proper performance of their particular duties. g. Violation of published University regulations, including regulations relating to entry and use of University facilities. h. Violation of published rules governing University residence halls.

Policy on Plagiarism

i. Deliberate alteration or misappropriation of computer records, data, software, etc. of the University or member of the University community. j. Breach or misuse of the Code of Computing Practice for the University of Windsor Computer Centre user.

Sanctions Defined a. Admonition. Notice to the student, orally or in writing, that s/he has violated student 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 severe disciplinary action. b. Censure. Written reprimand for violation of a specified regulation, including the possibility of more severe disciplinary sanction in the event of conviction for the violation of any University regulation within a period of time stated in the letter of reprimand. c. Disciplinary Probation. Exclusion from participation in privileges or extracurricular University activities as set forth in the notice of disciplinary probation for a specified period of time. d. Restitution. Reimbursement for damage or misappropriation of property. Reimbursement may take the form of appropriate service to repair or otherwise compensate for damages. e. Suspension. Exclusion from classes and other privileges or activities as set forth in the notice of suspension for a definite period of time. f. Expulsion. Termination of student status for an indefinite period. The conditions of readmission, if any is permitted, shall be stated in the order of expulsion. g. Exclusion from Campus. Denial of access to the campus for an indefinite period for non-academic misconduct. The conditions for removing this ban, if any, shall be included in the exclusion order.

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

S T U D E N T C O N T R A C T

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 laboratory 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: ______

GA’s Name: ______

Student’s Name (please print): ______

ID # ______Signature ______

Submit this copy to your GA.

Marking Scheme and Outline

MARKING SCHEME AND OUTLINE

All laboratory reports are marked out of 20.

Abstract (3 marks) The abstract is a crucial part of your laboratory report. It is a clear written paragraph which briefly conveys the intent of the laboratory exercise, mentions the methodology utilized during the laboratory period, summarizes important experimental observations, findings, and conclusions. 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 laboratory manual”, quoting the relevant pages. However, if there is a major deviation from experimental procedure or significant problem with the data arising from experimental error, you must state these differences in this section, as these deviations could influence the marks you receive in later sections (i.e., see Results and Calculations below).

Results and Calculations (Experiments 1-4: 8 marks & Experiment 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 your experimental data with literature values where appropriate. Show a sample calculation for each type of calculation required, including the equation used, the numbers substituted into the equation, unit analysis, and the final answer.

Discussion (Experiments 1-4: 7 marks & Experiment 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 underlying theory of the experiment. Important results should be compared with those found in the literature, and reasons for differences between experimental and literature data should be discussed. When there are several sources of possible error, indicate to what extent each may have affected the final results. You may also choose to include in your discussion possible applications of the experiment, suggested improvements or extensions to the experiment, and/or alternative methods of measurement. Answer the laboratory 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 laboratory 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 °C.3

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

Marking Scheme and Outline

Final Laboratory Grades

An attempt will be made to post unofficial laboratory grades prior to the final exam. It will be the responsibility of students to verify the grades posted as the marks one received for their laboratory write-ups. Students will have exactly one week after the unofficial laboratory grades are posted to dispute any marks they may have received. Students must bring in their marked laboratory report to the Laboratory Coordinator (Essex Hall, Room 175) to verify the mark received does not correspond with what was posted. Once the one week deadline has passed, the final laboratory grades will become official and will be submitted to the professor for final course grades. Any discrepancies, at this point, must be appealed. Please refer to the Registrar’s Office for details.

Marking Scheme and Outline

GENERAL LABORATORY OUTLINE for 59-240

1. Laboratory attendance is compulsory. 2. Laboratory reports must be turned in exactly one week after the laboratory. If your report is one day late, there is a 50% mark deduction. Reports more than two days late will not be accepted, and a grade of zero will be assigned. 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 this course). 4. Laboratories which are not completed receive a grade of zero. A completed laboratory refers to both attendance and performance in the laboratory, as well as handing in a completed report. 5. If the laboratory is not completed due to illness or family tragedy, the student may have the Student Medical Certificate (see next page) to the upper year Laboratory Coordinator in order to be 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 laboratories are offered. 7. If there is any evidence of plagiarism or duplication in laboratory reports among laboratory partners, friends or other students, all parties will receive a grade of zero and be subject to academic discipline, as per the University of Windsor Undergraduate Calendar.

Pre-Laboratory Preparation Read your laboratory manual carefully. You must be prepared to answer questions asked by your GA about the experiment and for surprise quizzes. You must know, or have written down, the physical constants of the chemicals you will be using. MSDS sheets are readily available for those who wish to consult them. You must also know how to prepare solutions relevant to your experiment (e.g. 4 M HCl from concentrated HCl) before you come to the laboratory. If there is any procedure, or equipment operation, with which you are not familiar, PLEASE ASK YOUR GA FOR HELP. This will help to prevent accidents and waste of valuable time.

DISPUTES, COMPLAINTS, AND POLICY Any disputes or complaints arising from the laboratory course should, in the first instance, be drawn to the attention of the GA. Should his/her decision be deemed unjust, representations may be made, in turn, to the Professor instructing the course, the Head of the Department of Chemistry and Biochemistry, and the Dean of Students.

ACKNOWLEDGEMENTS The current edition of this manual was prepared by Robert W. Schurko. Previous editions of this manual, as well as development and refinement of the experiments are credited to Paul Goulet, Cory Widdifield and Robert W. Schurko.

Student Medical Certificate

Experiment 1 – Determination of ΔcH: Bomb Calorimetry

EXPERIMENT 1

Determination of ΔcH: Bomb Calorimetry

Introduction

The Bomb Calorimeter

The calorimeter utilized in this laboratory experiment is composed of several parts and you must be familiar with many of them prior to starting this experiment (see Figure 1.1).

1 – High–precision Thermometer 2 – Thermometer Bracket 3 – Thermometer Support Bracket 4 – Reading Lens 5 – Thermometer Support Rod 6 – Motor 7 – Motor Pulley 8 – Stirrer Drive Belt 9 – Stirrer Pulley 10 – Stirrer Bearing Assembly 11 – Ignition Wire 12 – Stirrer Shaft with Propeller 13 – Bucket 14 – Calorimeter Jacket and Cover 15 – Oxygen Combustion Bomb

Figure 1.1

In addition to these items, one must also recognize the oxygen filling connection and the ignition unit. Beyond this (i.e., much of the fine structure of the oxygen combustion bomb), a GA will help you to understand how the different parts work and will also lead you through some of the important first steps (see Procedure).

Experiment 1 – Determination of ΔcH: Bomb Calorimetry

Measuring the Heat Evolved in a Combustion Reaction and Relation to ΔcH

The change in enthalpy associated with burning a sample is defined as its enthalpy of -1 combustion, ΔcH, and is expressed in units of kilojoules per gram (kJ g ). Oftentimes, a sample of benzoic acid is used to standardize a given bomb calorimeter, as its enthalpy of combustion is well-defined. When benzoic acid combusts, it can be described by the following chemical equation: 15 C H CO H  O  7 CO  3H O [1] 6 5 2 (s) 2 2 2(g) 2 (l)

For the following discussion, the system is defined as anything that is chemically involved in the combustion of the sample. According to the first law of thermodynamics, the change in internal energy of a system, ΔU, is related to the work done on the system, w, and the heat energy transferred to the system, q, and may be expressed algebraically as ΔU = w + q. Using temperature and volume as independent system variables, small changes in internal energy can be expressed as:  U   U  dU    dT    dV [2]  T V  V V If it is assumed that (i) no work is done on the system and (ii) the system will always occupy the same volume, then w = 0 and [2] simplifies accordingly:

 U  dU  dq    dT  CV dT [3]  T V By experimental design, the measurement of the change in internal energy of the system occurs indirectly, by determining the temperature change of the surroundings. This fact becomes important when considering the sign (i.e., ±) associated with the internal energy change of the system.

Enthalpy is defined as H = U + pV. Small changes in enthalpy can be represented algebraically:

dH  dU  d(pV)  dw  dq  d(pV)  CV dT  RTi dn [4] Where dn represents the change in the number of moles of gaseous species in the system. Here, it is assumed that the gaseous components of the system may be described by the ideal gas law. Note that only the change in internal energy is measured experimentally and a small correction (the RTi dn term) is applied to find ΔcH. Normally, the difference between the changes in the two state functions is very small and may be neglected.

Thus, if all of the heat evolved in the combustion reaction can be measured and if all the work done upon the surroundings is assumed to be negligible, then the internal energy change of the system due to combustion can be measured:

Δ H Δ U C(T T ) C T c  c   f  i    [5] Where C is understood to be the heat capacity of the calorimeter, in units of kJ K-1 (as the surroundings are not in the gas phase, the V subscript may be neglected). This heat capacity will be determined in the first part of the laboratory experiment by burning benzoic acid.

Experiment 1 – Determination of ΔcH: Bomb Calorimetry

Refining the Model

If ΔcH is to be determined in a fairly accurate manner, the model proposed above must be refined in order to account for additional factors.

As one may suspect, the calorimeter is not truly adiabatic and work is done upon the surroundings in the form of mechanical stirring. In an effort to account for these items, a radiation correction to the observed temperature change is made, and is represented by the following equation:

ΔT T T r (t t ) r (t t )  f  i  α 60  i  β f  60 [6]

Where Ti /ti and Tf /tf are the corrected temperatures (C) and the times (min) associated with the system at ignition and at the start of equilibrium, respectively, rα and rβ are the rates of -1 temperature change (C min ) before ignition and during equilibrium, and where t60 is the time at which 60% of the overall temperature change has occurred. The state of equilibrium is said to occur when, after ignition, the change in the observed temperature becomes constant for a period of 5 minutes. Corrected temperature values are to be applied according to the table below:

Table 1.1.

Range (C) Correction (C) Range (C) Correction (C)

19.0 – 20.3 – 0.013 24.9 – 25.3 – 0.007

20.4 – 22.7 – 0.012 25.4 – 25.7 – 0.006

22.8 – 23.3 – 0.011 25.8 – 26.0 – 0.005

23.4 – 23.8 – 0.010 26.1 – 26.3 – 0.004

23.9 – 24.3 – 0.009 26.4 – 26.6 – 0.003

24.4 – 24.8 – 0.008 26.7 – 27.8 – 0.002

Although the thermometer can read temperatures up to 35.000 C, it is not expected that readings above those listed will be made. If required, a GA will provide you with the corrections outside this range.

The final correction that must be done concerns the ignition wire. In order to ignite the sample, a small amount of nickel/chromium fuse wire is used and is partially consumed. The correction that is to be applied can be found simply by collecting all remaining portions of the fuse wire and measuring their length. For every 1.0 cm of fuse wire burned, a correction of 9.6 joules is applied (as an exercise, determine if this correction is to be positive or negative and make sure to point out your reasoning in your discussion of results).

Experiment 1 – Determination of ΔcH: Bomb Calorimetry

Materials

· 2000 mL volumetric flask · 2 – benzoic acid pellets · oxygen tank · 2 – unknown sample pellets · Parr oxygen bomb calorimeter, ignition · distilled water unit, oxygen filling adaptor, · nickel/chromium wire & accessories · wire cutters · cloth (for drying items)

Procedure

Parts A and B are each to be conducted in duplicate. If your results differ by more than 5 %, a third experiment should be conducted, time permitting.

A. Bomb Standardization

1. The GA will provide you with two benzoic acid pellets (0.7 – 1.0 g) (or you will help prepare pellets) and introduce you to some important aspects of the bomb calorimeter.

2. Determine the mass of a sample pellet and then transfer it to a fuel capsule.

3. Place the oxygen bomb head on the bomb head support stand, then put the fuel capsule with the sample in the electrode loop.

4. Cut about 10 cm of nickel/chromium wire and affix to the bomb head by placing the ends of the wire through the eyelets in each electrode (see right). Make sure that (i) the wire touches the top of the sample pellet and (ii) does not touch the sides of the fuel capsule.

5. Add approximately 1 mL of distilled water to the bottom of the bomb and fill the 2000 mL volumetric flask with distilled water.

6. Open the gas valve on the top of the bomb head, carefully transfer the bomb head into the oxygen combustion bomb, screw on the lid until it is hand tight, then close the gas valve.

7. Fill the bomb with about 25 atm of O2 (ask the GA to show you how to do this the first time). Slowly purge the O2 and refill the bomb to a pressure of about 25 atm once more.

8. Place the bomb into the calorimeter bucket and connect the ignition wires to the terminal sockets.

9. Slowly fill the calorimeter bucket, using all of the distilled water in the volumetric flask. Once done, place the lid on the calorimeter and attach the drive belt.

10. Let the system stand for about 2 minutes, then turn on the stirring motor and begin to take temperature readings (see Data Sheet).

11. Attach the unconnected lead wire to the common terminal (the other should already be in the 10 cm binding post of the ignition unit). 36

Experiment 1 – Determination of ΔcH: Bomb Calorimetry

12. When t = 3.0 min, depress the black button on the ignition unit for 3 – 4 seconds. The red light should come on for about half a second.

13. Continue to take measurements as outlined on the data sheet until the equilibrium period has been established (i.e., you can stop once the temperature change becomes constant over a period of 5 minutes).

14. Disconnect the lead wire from the common terminal, take off the calorimeter lid, wipe the end of the thermometer with the cloth provided and place lid onto appropriate stand.

15. Using the lifting handle, take the bomb partially out of the bucket (enough to expose the outsides of the terminal sockets to air), disconnect the ignition wires from the terminal sockets, and place the bomb on the counter top.

16. Without directly touching the bucket (you can touch the handle, for anything else use a cloth), dump the water out of it, and dry it and everything else that is wet with a cloth.

17. Slowly vent the gases from the bomb, then disassemble it, placing the bomb head on the appropriate stand.

18. Measure the amount of unburned wire, dry the insides of the bomb and remove any metal oxide deposits from the bomb interior, electrodes and fuel capsule by gently snipping/rubbing them with the wire cutters.

Note: If the sample did not completely combust, the GA will instruct you on how to dispose of the remainder.

B. Determination of Unknown Sample

1. The GA will provide you with one unknown sample, which you will press into two pellets (record the code number). Repeat the steps above using the unknown sample.

Results

1. Complete the data sheets on pages 39 – 40.

2. Create plots of temperature versus time and determine the parameters in [6].

3. Determine the corrected values for Ti and Tf according to Table 1.1.

4. Calculate and then tabulate: ΔT and the correction for consumed fuse wire for all trials, C

for part A, and ΔcH for part B. You may assume that ΔcU  ΔcH.

5. Determine average values for C and ΔcH. For each, calculate the percent difference between your two best readings.

6. Determine the identity of the unknown sample (see Table 1.2). Determine the percent error of your experiment.

Experiment 1 – Determination of ΔcH: Bomb Calorimetry

Table 1.2.

-1 -1 Compound Formula M (g mol ) ΔcH (kJ g )

Benzoic Acid C6H5COOH 122.12 – 26.43

Salicylic Acid C6H4(OH)COOH 138.12 – 21.90

Dibenzil (C6H5CH2)2 182.26 – 41.47

Hexamethylbenzene C6(CH3)6 162.27 – 43.95

Naphthoic Acid C11H8O2 172.18 – 29.68

Naphthalene C10H8 128.17 – 40.26

Lab Questions

1. After filling the bomb to a pressure of about 25 atm, why was the gas vented and refilled? What possible side reaction does this action help reduce? Explain.

2. What was the purpose of adding 1 mL of distilled water to the bottom of the bomb?

3. A student uses a bomb calorimeter and obtains the following data: Sample Run Compound Initial T (°C) Final T (°C) Weight Calibration benzoic acid (s) 0.218 g 23.84 24.97 Sample adipic acid (s) 0.326 g 24.38 25.61

(a) Determine the specific heat of the calorimeter. (b) Calculate the molar enthalpy of combustion of adipic acid, and compare (% difference) to the most recent literature values.

References

1. Atkins, Peter. Physical Chemistry, 7th edition Chapters 2 and 3 (or 8th edition Chapter 2). 2. Shoemaker, David P. and Garland, Carl W. (1970). Experiments in Physical Chemistry. (McGraw Hill). 3. The Parr Instrument Company, Manuals 204M, 205M, and 207M 4. NIST Chemistry WebBook. http://webbook.nist.gov/chemistry/

Recommended Reading

Small sections on calorimetry and thermochemistry in Lectures 6-8; corresponding sections in Atkins. See the back of the book for useful tables of thermodynamic constants.

Experiment 1 – Determination of ΔcH: Bomb Calorimetry

Data Sheets

A. Mass of benzoic acid pellet: Trial #1 ______g Trial #2 ______g Trial #3 ______g

Trial #1 Trial #2 Trial #3

t (min) T (C) t (min) T (C) t (min) T (C) t (min) T (C) t (min) T (C) t (min) T (C)

0 6.5 0 6.5 0 6.5

1 7 1 7 1 7

2 7.5 2 7.5 2 7.5

3 8 3 8 3 8

3.25 8.5 3.25 8.5 3.25 8.5

3.5 9 3.5 9 3.5 9

3.75 9.5 3.75 9.5 3.75 9.5

4 10 4 10 4 10

4.25 10.5 4.25 10.5 4.25 10.5

4.5 11 4.5 11 4.5 11

4.75 12 4.75 12 4.75 12

5 13 5 13 5 13

5.5 14 5.5 14 5.5 14

6 15 6 15 6 15

Experiment 1 – Determination of ΔcH: Bomb Calorimetry

Data Sheets

B. Mass of unknown sample pellet: Trial #1 ______g Trial #2 ______g Trial #3 ______g

Trial #1 Trial #2 Trial #3

t (min) T (C) t (min) T (C) t (min) T (C) t (min) T (C) t (min) T (C) t (min) T (C)

0 6.5 0 6.5 0 6.5

1 7 1 7 1 7

2 7.5 2 7.5 2 7.5

3 8 3 8 3 8

3.25 8.5 3.25 8.5 3.25 8.5

3.5 9 3.5 9 3.5 9

3.75 9.5 3.75 9.5 3.75 9.5

4 10 4 10 4 10

4.25 10.5 4.25 10.5 4.25 10.5

4.5 11 4.5 11 4.5 11

4.75 12 4.75 12 4.75 12

5 13 5 13 5 13

5.5 14 5.5 14 5.5 14

6 15 6 15 6 15

Experiment 2 – Vapour Pressure of Pure Liquids

EXPERIMENT 2

Vapour Pressure of Pure Liquids

Introduction

In this experiment, the relationship between the vapour pressure of pure liquids (methanol and ethanol) and temperature is studied. From careful experiments measuring the vapour pressure with respect to temperature, the molar enthalpies of vaporization (in kJ mol-1) can be determined for both methanol and ethanol.

When a pure liquid is placed in an evacuated bulb, molecules leave the liquid phase and enter the gas phase until the vapour pressure in the bulb reaches a definite value. This is determined by the nature of the liquid and its temperature. This pressure is called the vapour pressure of the liquid at a given temperature. The vapour pressure of a pure liquid is independent of the quantity of liquid and vapour present, as long as both phases exist in equilibrium with each other at the specified temperature. If the temperature increases, the vapour pressure also increases up to the critical point, where the two phases become a single homogeneous, one-phase supercritical fluid.

If the pressure above the liquid is maintained at a fixed value, then the liquid may be heated up to a temperature at which the vapour pressure is equal to the external pressure. At this point, vaporization will occur by the formation of bubbles in the interior of the liquid as well as at the surface. This is the boiling point of the liquid at the specified external pressure. Clearly the temperature of the boiling point is a function of the external pressure; in fact, the variation of the boiling point with external pressure is seen to be identical with the variation of the vapour pressure with temperature.

It can be shown that a definite relationship exists between the values of pressure, p, and the temperature, T, for a pure liquid and its vapour at equilibrium: dp ΔS dT = ΔV [1] where: · dp and dT refer to infinitesimal changes in pressure and temperature for a pure substance with both phases present in equilibrium. · ΔS and ΔV refer to the change in entropy, S, and volume, V, when one phase transforms to the other at constant pressure and temperature.

Since the change in state is isothermal, and ΔG is zero, ΔH - TΔS = 0. ΔS may be replaced by ΔH/T, giving: dp ΔH dT = TΔV [2] which is the Clapeyron equation. It is an exact expression which may be applied to phase equilibria of all kinds, although it is presented here in terms of a liquid-vapour phase transition of a pure substance (i.e., one component).

Experiment 2 – Vapour Pressure of Pure Liquids

Since the enthalpy of vaporization, ΔvapH, is positive, and ΔV is positive for vaporization, the vapour pressure must increase with increasing temperature.

For vapour-liquid phase transition equilibria in the range of vapour pressures less than 1 atm, one may make the following assumptions:

1. The molar volume of the liquid is negligible in comparison to that of the vapour, so that ΔV = Vg, where Vg is the volume of the vapour. Therefore,

dp ΔvapH = [3] dT TVg Since d ln p = dp/p, and d(1/T) = – dT/T 2, [1] can be rewritten in the form

dlnp ΔvapH RT ΔvapH = - = - [4] d(1/T) R pVg ZR where Z is the compressibility factor. The compressibility factor takes into account real gas behaviour, thereby yielding increasingly accurate results for determining the change in temperature with pressure. The compressibility factor for the vapour is equal to:

pVg Z = RT [5] The compressibility factor may also be calculated using the Berthelot equation: 2 9pT  6T  c  c Z = 1 + 201 - 2 [6] 128pc T  T 

where: Tc = critical temperature pc = critical pressure

2. [4] is a convenient form of the Clapeyron equation. If the vapour is a perfect gas (Z = 1) and ΔvapH is independent of the temperature, then a plot of ln p vs. 1/T yields a straight line, the slope of which is equal to ΔvapH. For many liquids, ln p is essentially a linear function of 1/T, which implies at that ΔvapH/Z is almost constant.

The Clausius-Clapeyron equation, which is derived from the Clapeyron equation, relates the enthalpy of vaporization, ΔvapH, of a pure liquid to its vapour pressure at various temperatures:

ΔvapH 1 ln p = - R T + C [7] where: ln p = the natural logarithm of the vapour pressure, ΔvapH = the heat of vaporization, R = the gas constant, T = temperature (in K), and C = a positive constant. This equation resembles a linear equation of the form: y = ax + b.

Experiment 2 – Vapour Pressure of Pure Liquids

Materials

· rubber stopper assembly · 2 – plastic caps · tubing with two connectors · gas chromatography septa · 10 mL syringe · Vernier Lab Pro · 2 – 100 mL dual neck round bottom flasks · Vernier gas pressure sensor · 4 – 600 mL beakers · Vernier temperature probe · 2 – thermometers (for water baths) · methanol · 15 mm washers · ethanol absolute · 3 – 100 mL beakers (1 for waste)

Procedure

A. Preparing for Data Collection

Pressure and temperature will be measured using a pressure sensor and a temperature probe.

1. Use the 600 mL beakers to prepare four water baths, one in each of the following temperature ranges: 0 to 5 C, 10 to 15 C, 20 to 25 C (use room temperature water), and 30 to 35 C (“hot” water can be obtained directly from the tap).

2. Prepare the temperature probe and pressure sensor for data collection:  Plug the temperature probe into Channel 2 of the Lab Pro.  Plug the pressure sensor into Channel 1 of the Lab Pro.  Assemble the apparatus as per Figure 2.1. Ensure that the stopcock is open, the stopper is tightly inserted into the round bottom flask, and plastic cap and septa are tightly screwed in place. Do not “Hulk Hogan” it! (i.e., use gentle force)

Figure 2.1

Experiment 2 – Vapour Pressure of Pure Liquids

3. Prepare the computer for data collection by opening the Chemistry with Computers folder. Then open the file: Vapour Pressure (Exp. 10. Gas Pressure-Stainles.MBL). The vertical axis will have pressure scaled from 90 to 135 kPa. The horizontal axis will have temperature scaled from 0 to 50 C.

4. Close the stopcock and immerse the round bottom flask into the room temperature (20 – 25 C) water bath.

Note: Do not apply pressure to the rubber stopper, since this will alter the pressure inside the round bottom flask, and ruin your results!

5. After 30 seconds or so, click on the Logger Pro software. This will provide a real time display of the temperature and pressure readings. When equilibrium has been reached (i.e., stable pressure and temperature readings), click “KEEP” on the Logger Pro software. The first pressure/temperature data pair is now stored.

6. Replace the room temperature water bath with the high temperature water bath (30-35 °C). Observe the pressure reading of the flask. The pressure reading of the flask should increase by 2 – 5 kPa. If no pressure increase is observed or the pressure goes up and then slowly decreases, then check all of the seals for leaks. If there are no leaks, proceed to Part B.

B. Measurement of vapour pressure of methanol: collection of data at room temperature

Starting with a water bath at (20 – 25 C):

1. Place the flask and temperature probe in the room temperature water bath. Once the pressure has stabilized, click “KEEP” and record the p and T data pair for the empty flask on your data sheet (round to the nearest 0.1 kPa).

2. Close the stopcock and immerse the flask into the water bath, with the entire flask covered as shown in Figure 2.1.

Note: In order to obtain accurate results, take the same care as you did in part A in ensuring the caps and stopper are tightly sealed.

3. Pour about 10-15 mL of methanol into a 100 mL beaker. Draw 3 mL of methanol into your syringe. Briskly wipe off the syringe needle with a Kimwipe.

Note: Consult the GA on the proper use of a syringe – it is very important that air bubbles be removed and that you have the proper volume in the syringe.

Experiment 2 – Vapour Pressure of Pure Liquids

4. Gently insert the syringe needle through the septum and inject the 3 mL of methanol into the round bottom flask.

5. Return the plunger of the syringe back to the 3 mL mark, ensuring that you do not pass this mark. Be careful here!

6. Gently remove the syringe from the septum.

7. Monitor the pressure and temperature data. When the p and T readings stabilize, meaning that the equilibrium between methanol liquid and vapour has been established, click “KEEP”. Record the pressure/temperature pair on your data sheet. Click on “KEEP” multiple times to record a series of data points from which you can calculate an average.

C. Measurement of vapour pressure of methanol: collection of data at other temperatures

Obtain pressure and temperature data in a similar manner by replacing the room temperature water bath with the 0 to 5 C, 10 to 15 C and 30 to 35 C water baths. Pressure/temperature data pairs for all of these trials should be recorded on your data sheet.

Note: Be sure to wait for pressure/temperature equilibration.

D. Measurement of vapour pressure of ethanol

Dispose of the methanol as instructed by the GA. Thoroughly clean and dry all of the glassware and the syringe. Repeat the above steps using ethanol, and water baths with temperature ranges of 0 to 5 C, 10 to 15 C, 20 to 25 C, and 30 to 35 C.

Results

1. Fill out data sheet on the next page.

2. To obtain the vapour pressures of methanol and ethanol, air pressures must be subtracted from each of the measured pressure values. These corrected air pressures are determined with the following relationship derived from the perfect gas law:

p2 p1 = T2 T1

· p1 is the atmospheric pressure (in kPa) of the empty flask and T1 is the temperature (in Kelvin) of the room temperature water bath (part A).

· T2 is the temperature of the water bath for any other measurements.

· p2 is the corrected air pressure, which is what you will solve for.

Experiment 2 – Vapour Pressure of Pure Liquids

3. Obtain the vapour pressures by subtracting the corrected air pressures from the measured pressures.

4. For both methanol and ethanol, use Excel (or other appropriate spreadsheet program) to plot graphs of vapour pressure as a function of temperature (in Kelvin) from the four data pairs you collected.

5. Using Excel, plot ln p as a function of l/T for both methanol and ethanol. In each case, determine the enthalpy of vaporization from the slope, – ΔHvap/R.

Lab Questions

1. (a) Isopropanol boils at 82.5 C and 1 atm. Estimate the heat of vaporization using Trouton’s rule:  1 1 Δ vapHT,m  (85 J K mol )Tbp and compare this value to the literature value for isopropanol (% difference). Carry out the same calculation for carbon tetrachloride, which boils at 76.7 °C. (b) Why is it reasonable to estimate the enthalpy of vaporization for carbon tetrachloride in this manner, but not for isopropanol?

2. At what temperature will isopropanol and carbon tetrachloride have the same vapour pressure?

References

1. Atkins, Peter. Physical Chemistry. 7th edition. Chapter 6, Sections 6.1 and 6.2 or 8th edition Chapter 4, Sections 4.1 and 4.2.

2. Shoemaker, David P. and Garland, Carl W. (1970). Experiments in Physical Chemistry. (McGraw Hill).

3. Thomson, G. W. Chem. Rev., 38, 1946: 1.

Recommended Reading

Compressibility factor, Z: Lecture 4; Critical pressures and temperatures: Lecture 4, end: see critical constants and principle of corresponding states at end of lecture; Trouton’s Rule: Lecture 11; Clausius-Clapeyron equation: Lecture 14, end (Liquid-Vapour boundaries)

Experiment 2 – Vapour Pressure of Pure Liquids

Data Sheet

Empty Flask:

Pressure: ______

Temperature: ______

Trial #1 (room temp.) #2 (0 C to 5 C) #3 (10 C to 15 C) #4 (30 C to 35 C)

Substance CH3OH C2H5OH CH3OH C2H5OH CH3OH C2H5OH CH3OH C2H5OH

Temp. (C)

Temp. (K) Measured pressure (kPa) Corrected air pressure (kPa) Vapour pressure (kPa)

Experiment 2 – Vapour Pressure of Pure Liquids

Experiment 3 –Surface Tension of n-Butanol and Amount Adsorbed

EXPERIMENT 3

Surface Tension of n-Butanol and Amount Adsorbed

Introduction

Surface Tension

A surface is an interface between two phases, and may be one or more molecular layers thick. In principle, concepts of equilibrium concentrations and thermodynamic properties can be applied to treat surfaces in the same general way as all other phases of matter.

Within a liquid, all molecules experience very similar intermolecular interactions, being pushed and pulled with equal force from all sides. However, at the gas-liquid interface or surface of a liquid, the molecules do not experience the same forces as those “buried” below the surface. These surface molecules experience only the forces from molecules within the bulk phase and essentially nothing from the other side of the interface; correspondingly, these molecules have a higher chemical potential (μ) than those below the surface. Therefore, increasing the surface area of a liquid requires an input of energy, and liquids have a tendency to minimize their surface areas by retaining the maximum possible number of molecules in the bulk volume. This gives rise to a unique property of liquid–gas interfaces, known as surface tension, which may be measured using a variety of different techniques. Surface tension, γ, is commonly expressed as a force per unit distance (e.g., N m-1) or an energy per unit area (e.g., J m-2). This corresponds to the minimum work (reversible work) needed to increase the surface by one unit of area.

Let us consider a solution with N moles of solute. The solute exists in two regions in the solvent: (i) in the bulk “interior” of the solvent and (ii) at the surface or “interface” of the solvent. The solution has a uniform concentration right up to the surface region, with bulk concentration NV/V = c (units of moles per unit volume). In the surface region, there is a solution with a slightly higher concentration (i.e., excess of solute compared to the bulk phase), with a surface concentration NA/A = u (units of moles per unit area).

Any system at constant T and p seeks a state of lowest free energy, and the maximum work done by the system is equal to the free energy decrease. In the case of solute/solvent mixtures, equilibrium conditions are reached when the free-energy decrease due to lowering surface tension is balanced by an opposing tendency for free-energy increase due to increasing non-uniformity of the solute near the surface.

Imagine a solution of volume (V), surface area (A), bulk concentration (c), surface concentration (u), surface tension (γ), and bulk osmotic pressure (Π), all at constant p and T. For arbitrary changes in the area and volume of the solution (i.e., dA and dV), the free-energy change can be written as dG = γdA - ΠdV [1]

This is an exact differential (see Further Information 1 in Atkins, Physical Chemistry, 6th or 7th Edition). Thus, from the reciprocity relation, we may write

Experiment 3 –Surface Tension of n-Butanol and Amount Adsorbed

 γ   Π        [2]  V A  A V

Since at constant p and T, the surface tension and osmotic pressure are completely determined by the bulk concentration, we can rewrite [2] as

γ  c  Π  c        [3] c  V A c  A V

If the total number of moles of solute is N = cV + uA, then

N  uA c  [4] V This can be differentiated with respect to volume and area to yield:  c  c  c  u        [5]  V A V  A V V As well, for a perfect solute, the osmotic pressure can be calculated from N Π  RT  cRT [6] V which can be differentiated with respect to concentration to yield dΠ  RT [7] dc If we combine [3], [5] and [7], we obtain the following expression: u 1 dγ   [8] c RT dc which is known as the Gibbs isotherm. It can also be recast in the more useful forms below: 1 dγ dγ N RT u   or  uRT   A [9] RT d (lnc) d (lnc) A

This derivation implies that the surface tension of a solution differs from that of the pure solvent because of the adsorption of solute. If there is no adsorption of solute, γ doesn’t change. If there is adsorption of excess solute at the surface (solute attracted to the surface), γ decreases as the concentration c of solute in the solution is increased. This is a well known behaviour of a soap solution which has lower surface tension than pure water. If the solute avoids the surface region (i.e., “negative adsorption”), γ will increase with increasing c, and the solution will have a greater surface tension than the solvent. This commonly happens with ionic solutes in water which tend to decrease foam formation (for example, it is difficult to get soap suds in salt water or “hard” water containing Ca(HCO3)2).

Experiment 3 –Surface Tension of n-Butanol and Amount Adsorbed

If we plot surface tension against ln c, we can observe the following: 1. If positive adsorption occurs the curve should have a negative slope.

2. If the plot is linear, the derivative on the left side of [9] is constant, independent of c. The amount of solute adsorbed per unit area must be constant, and while the concentration of solute in the solution is changing, the concentration in the surface region is not. This likely means that adsorption saturation has been reached, producing a complete monolayer of adsorbed solute on the surface. From the slope of a plot of γ vs. ln c, we can find NA/A and hence the number of molecules of solute per unit area in a complete monolayer.

All this can be done from surface tension measurements, with the aid of the second law of thermodynamics, and without any direct determination of the amount adsorbed.

Determination of Capillary Diameter

The capillary rise method is used to study the change in surface tension as a function of concentration for aqueous solutions of n-butanol. The objective is to show that the solute forms a complete monolayer at the surface and find the area occupied by a molecule in that monolayer, using the Gibbs adsorption equation.

It is known that in the absence of external forces, a body of liquid tends to assume a shape of minimum area. When a liquid is in contact with a solid surface, a specific surface free energy exists for the interface, or interfacial tension, γ1,2. A solid surface itself has a surface tension, γ2, which is often large in comparison with the surface tensions of liquids.

Suppose a liquid with surface tension, γ1, is in contact with a solid with surface tension, γ2. Then there is an interfacial liquid-solid surface tension, γ1,2. Under what circumstances will a liquid film freely spread over the solid surface and “wet” it? Wetting will happen if the free energy of the entire system decreases as the result of creating a liquid-solid interface, such that

γ  γ  γ  0 [10] 1 1,2 2

If a capillary tube of radius, r, is dipped into a liquid with a surface tension γ and density ρ, and if the liquid wets the walls of the capillary so that it adheres to the walls with contact angle θ = 0, then the height, h, of liquid in the capillary tube (measured as in Figure 3.1) is given by balancing the upward force of surface tension against the downward force of the weight of the column of liquid. This will give the equation, 1  r  γ  h  rρg [11] 1   2  3  where r/3 is a correction for the amount of liquid below the meniscus. However, if [10] is not obeyed, then we have γ1 cosθ  γ1,2  γ2  0 [12]

Experiment 3 –Surface Tension of n-Butanol and Amount Adsorbed

For some value of θ, the liquid will not tend to spread indefinitely on the solid surface but will tend instead to give a “contact angle, θ”. This may be the case with aqueous solutions or water on glass surfaces that are not entirely clean. [11] can then be rewritten as:

1  r  γ cosθ  h  rρg [13] 1 2  3 

In practice, the contact angle finally attained is somewhat variable, depending on whether the liquid is advancing over the solid surface or receding from it. If the same height is obtained regardless of whether the liquid was allowed to rise from below or fall from above in the capillary, it may be assumed that [12] and [13] are almost certainly valid.

Materials

· 4 – 100 mL bottles · graduated capillary tube · 200 mL volumetric flask · 50 mL volumetric pipette · 100 mL graduated cylinder · 10 mL graduated pipette · 100 mL beaker · 2 pipette bulbs · large test tube · tygon tubing (attached to vacuum) · 2-holed stopper · tweezers · thermometer · n-butanol

Procedure

Assemble the apparatus as shown in Figure 3.2. The capillary is stored in concentrated nitric acid. The capillary should be rinsed with water and dried with a vacuum hose and Kimwipes between measurements. Pour water into the large test tube to roughly match the scale shown in Figure 3.2. Adjust the capillary upward or downward until the outside liquid level is above the lowest graduation of the capillary.

Figure 3.1 Figure 3.2

Experiment 3 –Surface Tension of n-Butanol and Amount Adsorbed

1. The height, h, is the distance between the water levels inside and outside of the capillary tube. Determine h of the capillary rise for pure water at room temperature (take a minimum of 4 readings). Attach the bulb to the pressure inlet. Gently squeezing the bulb, force the height of the column of liquid to near the top of the capillary tube. Then, while the bulb is still compressed, remove it from the pressure inlet. The pressure release will allow the column of liquid to fall to its equilibrium position. If there is not good agreement among these readings, clean the capillary using concentrated HNO3 and repeat the measurements.

2. In the 200 mL volumetric flask, prepare an aqueous solution of 0.8 M n-butanol (ρ(n-butanol) = 0.810 g mL-1). This is the starting solution used for consecutive dilutions.

3. Using a 50 mL Pasteur pipette, withdraw 50 mL of solution from the volumetric flask, and empty into a 100 mL bottle. The 100 mL bottle now contains 50 mL of 0.8 M solution. Measure h for the 0.8 M solution in the bottle (as described in steps 1 and 2)

4. Add distilled water to the volumetric flask to bring the solution level up to the mark. This is now a 0.6 M solution of aqueous n-butanol. Again, withdraw 50 mL of the solution, and empty into a 100 mL bottle. Measure h for this solution as described in steps 1 and 2.

5. Repeat the above procedure of dilution and extraction to produce three more solutions, and measure h for each. Be sure to calculate the concentrations of these new solutions and record this data.

Calculations

1. From the data obtained for pure water, calculate the capillary radius, r, [11], then calculate the surface tension of each solution studied. For the n-butanol solution, one may assume that the density is equal to that of pure water. At 25 C for pure water the surface tension is 72 dyne cm-1 and the density is 0.9970 g cm-3.

2. Plot surface tension, γ, against the natural logarithm of the bulk molar concentration of n-butanol in the solution, c, and determine the slope.

3. Use the slope of the plot, with the aid of [9], to find the amount of n-butanol adsorbed in moles per unit area (NA/A). Convert this to molecules per square Ångstrom, then find the “effective cross-sectional area” per n-butanol molecule of absorbed butanol in Å2.

Lab Questions

1. Explain the difference between capillary rise and capillary depression.

2. Would the equation γ  gρrh(cos θ) be altered if the liquid were mercury, which does not wet glass? Explain. What would be the sign of h for mercury? Explain.

Experiment 3 –Surface Tension of n-Butanol and Amount Adsorbed

3. In an experiment to determine how well benzene wets glass, benzene was found to rise to 1.71 mm in a certain capillary tube at 20.0 C. In the same tube, water rose 4.90 mm at 20.0 C. Calculate the contact angle for benzene, given that at 20 C, ρbenzene = 0.879 g -3 -3 -2 -1 -2 -1 cm , ρwater = 0.998 g cm , γbenzene = 2.89  10 N m and γwater = 7.28  10 N m . The contact angle for water is assumed to be θ = 0°.

References

1. Atkins, Peter. Physical Chemistry. 7th edition, Chapter 6 or 8th edition Chapter 4.

2. Alberty, Silbey. (2001). Physical Chemistry. 3rd edition. Chapter 6, Section 6.4.

3. Shoemaker, David P. & Garland, Carl W. (1970) Experiments in Physical Chemistry.

Recommended Reading

Surface tension and capillary action: Lecture 15

See also: corresponding notes in Atkins.

Experiment 4 –Heat of Reaction in Solution: Constant Pressure Calorimeter

EXPERIMENT 4

Heat of Reaction in Solution: Constant Pressure Calorimeter

Introduction

In this experiment, a constant pressure calorimeter is used to determine the enthalpies of dissolution for several common salts. This is accomplished by monitoring changes in temperature while dissolution reactions are carried out in an adiabatic vessel. The energy change, q, for a reaction that takes place in a calorimeter, can be determined by multiplying the net temperature change, ΔT, by the heat capacity of the calorimeter and its contents, C. q = CΔT [1] The calorimeter constant, C, of a calorimeter and its contents, can be determined experimentally by standardization procedures, and must be known in order to calculate the energy change in the form of heat, q, for a given reaction. Three different procedures that are available for standardizing calorimeters are:

1. Chemical standardization using a “TRIS” (tris(hydroxymethyl)aminomethane) sample. This involves the precise and reproducible exothermic reaction of TRIS with 0.1 M HCl. 2. Comparison standardization which involves the use of samples whose enthalpy changes are known and whose thermochemical behaviour is similar to that of the unknown material. 3. Electrical standardization which requires an electric heating probe, a uniform power supply, a high precision voltmeter, and a precise interval timer.

The heat capacity of the constant pressure calorimeter used in this experiment (including its contents), can be determined by running several calibration trials where the calorimeter is operated in the usual manner, but where the reactants release a known amount of energy. For example, a very well controlled reaction of 0.5 g of TRIS, dissolved in 100 mL of 0.1 M HCl, will evolve 245.76 J g-1 of TRIS at 25 C. This standard value can be used to determine the heat capacity of the calorimeter and its contents (see Procedure, part A).

This experiment is performed at constant pressure. The heat of dissolution at constant pressure, qp, is equivalent to the enthalpy change, ΔdissHT, at the mean reaction temperature. Molar enthalpies of dissolution, ΔdissHT, can be obtained from the following relationship: q Δ H = p [2] diss T n

where: ΔdissHT = molar enthalpy of dissolution at mean reaction temperature, T; n = quantity of sample used (in moles); q = heat (i.e., calorimetric energy change).

Experiment 4 –Heat of Reaction in Solution: Constant Pressure Calorimeter

Materials

· solution calorimeter · scoopula · multimeter · mortar and pestle · 500 mL volumetric flask · concentrated HCl · 100 mL beaker · tris(hydroxymethyl)aminomethane (TRIS) · 100 mL graduated cylinder · potassium iodide · weighing dishes · potassium nitrate · 10 mL graduated pipette · potassium chloride · pipette bulb · ammonium nitrate

Procedure

The experimental procedure can be divided into two major sections. The first of these sections involves the determination of the heat capacity of the calorimeter and its contents at constant pressure by a calibration with TRIS/HCl. The second section employs the average heat capacity found in the initial standardization to determine the enthalpies of dissolution for three common salts in water (choose 3 of the 4 available salts).

Calibration and Use of Calorimeter (GA will demonstrate this - so read carefully!)

1. The multimeter should be properly connected to the constant pressure calorimeter, and should be set to read voltage (i.e., ).

2. When calorimeter is in “zero” mode, the multimeter should read 0.000 V. In the “null” mode it should also read 0.000 V. In the “cal” mode the calorimeter should be adjusted until the multimeter reads 1.000 V (knob on upper right hand side of calorimeter labelled “cal”). Finally, in the “read” mode, the temperature settings on the calorimeter should be adjusted to bring the voltage on the multimeter to 0.000 V. In this way, temperatures can be read directly from the calorimeter.

3. Once calibrated, the calorimeter will be left in “read” mode for the remainder of the experiment. All temperature measurements will be made by zeroing voltage, and then reading the calorimeter.

4. Before reading temperatures, always allow voltages to stabilize.

5. If the multimeter shuts itself off during the experiment (power saving function), simply switch it off, and then back on again. This will restore its function.

Part A: Determination of the Calorimeter Constant (C)

A sample of TRIS is dissolved in dilute HCl in a controlled reaction for which the amount of heat evolved is known. This standardization is performed in triplicate to obtain an average C.

1. Using an analytical balance, mass out exactly 100.00 ± 0.50 g of 0.100 M HCl (prepared by dilution from concentrated HCl) into a graduated cylinder (approximately 100 mL). Add this to the Dewar inside of the calorimeter.

Experiment 4 –Heat of Reaction in Solution: Constant Pressure Calorimeter

2. Weigh accurately 0.50 ± 0.01 g of TRIS in the 126C Teflon dish of the calorimeter (the salt should be ground into powder using a mortar and pestle to ensure complete reaction). Note the exact mass for use in subsequent calculations.

3. Assemble the rotating cell, and place it in the calorimeter. (This will be demonstrated by your GA, and additional instructions are included in the document “Calorimeter Operation” that is provided with the experimental materials.)

4. Allow the reactants to come to thermal equilibrium (stable voltage) in the closed calorimeter while stirring. Once equilibrium is established, record the initial temperature, Ti, and start the dissolution reaction (quickly push down rod without interrupting its rotation).

5. At this point an enthalpy change will take place in the calorimeter, and a temperature change will be observed. When the temperature stabilizes in the calorimeter, stop stirring, and record the final temperature, Tf.

6. The calorimeter can now be opened. All parts should be cleaned and dried to prepare for the next trial. If all of the solid has not been dissolved, the trial must be repeated.

7. Calculate the change in temperature, ΔT, for each trial and use it to calculate T, which is equal to the temperature at which 63% of the change has taken place.

T = 0.63ΔT + Ti [3] 8. Calculate the energy change for each trial using:

-1 -1 -1 q/(J) = mTRIS/(g)[245.76/(J g ) + 1.436/(J g °C )(25.00 °C - T)] [4]

where q = energy change in J m = mass of TRIS in grams T = calibrated temperature

Note: The term: 1.436(25.00 C - T) adjusts the heat of reaction to temperatures above or below the 25 C reference temperature.

9. Calculate the calorimeter constant for each trial using the equation: q C = ΔT [5]

10. Calculate an average calorimeter constant. This is the value that will be employed for the remainder of the experiment. (Note that we have assumed that the heat capacity of 100 g of a dilute aqueous solution is equal to that of 100 g of pure water. This is a valid assumption, since the specific heat capacity of 0.1 M HCl at 25 C is 4.180 J C-1 g-1 and the specific heat capacity of water is 4.184 J C-1 g-1).

Experiment 4 –Heat of Reaction in Solution: Constant Pressure Calorimeter

Part B: Determination of the Enthalpies of Dissolution for Salt Solutions

1. Determine the change in temperature, ΔT, for the dissolution of 0.50 ± 0.01 g of three different ground salts in 100.00 g of distilled H2O. Note the exact mass for each trial. These three salts will be assigned to you by your GA, and will be chosen from KI, KNO3, KCl, and NH4NO3. Perform two trials for each salt you are assigned using the same techniques as were used for the TRIS/HCl calibration.

2. Using [1] and [2] determine the energy change, q, for each trial. Calculate an average enthalpy of dissolution, ΔdissHT, for each of the salts.

3. Compare the values that you calculate with the enthalpies of dissolution of these salts (in kJ mol-1), with those found in the literature, and calculate a percent error for each.

Lab Questions (use data tables in the back of Atkins).

1. Calculate the enthalpy of neutralization ΔneutH° of KOH(aq) with an equimolar amount of HCl.

2. Demonstrate that for strong acids and bases that the enthalpy of neutralization can simply +  be determined by calculating the ΔrH° for H (aq) + OH (aq)  H2O(l). Why does this equation give the same result as question 1?

-1 3. The enthalpy of neutralization for the first proton of H2S is -33.7 kJ mol . Calculate the first acid ionization energy for H2S (aq).

 -1 2- 4. The enthalpy of neutralization of HS (aq) is -5.1 kJ mol . Calculate the ΔfH° of S (aq).

5. Calculate the ΔsolnH° of hydrogen chloride.

References

1. Atkins, Peter. (1998). Physical Chemistry. 7th edition or 8th edition, Chapter 2.

2. Parr Instrument Co. Solution Calorimeter Instruction manual.

3. Alberty, Silbey. (2001). Physical Chemistry. 3rd edition. Chapter 2, Section 2.15.

Recommended Reading

Thermochemistry: Lecture 8

Accompanying problems; see also: Atkins, Chapter 2.

Experiment 5 –Liquid-Vapour Equilibrium in a Binary System

EXPERIMENT 5

LIQUID-VAPOUR EQUILIBRIUM IN A BINARY SYSTEM

Introduction

In this experiment, the compositions of the vapour and liquid phases of several boiling mixtures of cyclohexane and acetone are determined by refractive index measurements. These compositions are then plotted against temperature to produce a liquid-vapour phase diagram for this binary system.

Distillation is a convenient technique for determining the liquid-vapour phase diagrams of binary liquid systems. When a homogeneous two-component liquid is distilled, the composition of the vapour is generally different from that of the liquid. The vapour pressures of the components of an ideal solution of two volatile liquids are related to the composition of the liquid mixture by Raoult’s Law: p  x p * A A A [1] pB  xB pB *

If we substitute these expressions into Dalton’s Law of Partial Pressures, p = pA + pB, we get the following expression for the total vapour pressure, p, of the mixture:

p  x p * x p * [2] A A B B

Given that xA + xB = 1, we get: p  p *  (p * - p *)x [3] B A B A where: · p is the total vapour pressure, * * · pA and pB are the vapour pressures of pure A and B, respectively, and

· xA and xB are the mole fractions of A and B.

Raoult’s Law is an excellent approximation for binary liquid-vapour systems when the mole fraction of one component is close to unity. Major deviations from Raoult’s Law occur when neither mole fraction is close to unity, or in considering the behaviour of the dilute component of the binary mixture. At constant temperature, if the vapour pressure of solution is higher than what is predicted by Raoult’s Law, the system exhibits a positive deviation. If the vapour pressure is lower, it is said to exhibit a negative deviation. The positive and negative deviations arise from homogeneous (i.e., A—A, B—B) and heterogeneous (i.e., A—B) molecular attractions, respectively. For example, a positive deviation implies that the homogeneous attractions are stronger than the heterogeneous ones. These deviations may be large enough to produce minima and maxima in vapour-pressure and boiling point curves, as shown in Figure 5.1.

Experiment 5 –Liquid-Vapour Equilibrium in a Binary System

Figure 5.1. Schematic vapour pressure and boiling point diagrams for systems showing (a) a strong positive deviation, and (b) a strong negative deviation from Raoult’s Law (from Shoemaker, Garland, and Nibler, 6th Ed.).

At the maxima or minima in these curves, the vapour and liquid compositions are the same, and there is a point of tangency between the curves L and V and L* and V* at the minima or maxima. Such systems exhibiting these minima and maxima are called azeotropes, and are of major importance in connection with distillation. The point of tangency in the temperature-composition diagrams is called the azeotropic temperature, and represents a constant pressure and temperature at which the two components of the binary mixture cannot be separated from one another by distillation. Binary mixtures exhibiting the so-called “positive deviation” on the pressure-composition graphs are referred to as low-boiling azeotropes, since only the liquid phase exists below the azeotropic temperature. Similarly, binary mixtures exhibiting the “negative deviation” are referred to as high-boiling azeotropes, since only the vapour phase exists above this temperature.

Materials

· Abbe refractometer and light source · 400 mL beaker · 2 – dual necked boiling flasks · rubber septa · condenser · 10 mL graduated pipette · 20 mL syringe · pipette bulb · Variac · 3 – 100 mL beakers · heating mantle · 2 – condenser hoses (Tygon) · 2 – thermometers · acetone · one-holed rubber stopper · cyclohexane

Experiment 5 –Liquid-Vapour Equilibrium in a Binary System

Procedure

A binary liquid-vapour phase diagram is to be constructed for the acetone-cyclohexane system by carrying out distillations of several different mixtures of the two components and determining their liquid (mixture remaining in the distillation flask) and vapour (distillate) compositions at their boiling points. The boiling point temperatures of the samples can be plotted against the compositions of the liquid and vapour phases (2 plots) in order to produce the phase diagram, and from this the azeotropic composition and temperature can be determined.

Distillation

1. Set up the distillation apparatus as shown in Figure 5.2 (Be sure to include grease and boiling chips in this step). The rubber stopper should be fit tightly into the boiling flask to prevent it from being dislodged during heating.

Figure 5.2.

2. Carry out distillations. This will involve 10 mixtures of different compositions being heated to boiling, and then sampled for refractive index measurements. The compositions of these samples are summarized in Table 5.1. To prepare the first 5 mixtures, simply add to the existing mixture each time (as per Table 5.1) without emptying the flask. When the first 5 mixtures have been distilled, the flask can then be emptied and washed to prepare for the next 5 distillations. These mixtures can also be prepared in an additive fashion (i.e., no washing in between samples).

Experiment 5 –Liquid-Vapour Equilibrium in a Binary System

Table 5.1

Composition

Mixtures 1 – 5 Mixtures 6 – 10

1 25.0 mL cyclohexane 6 25.0 mL acetone

2 add 2.5 mL acetone 7 add 1.5 mL cyclohexane

3 add 3.0 mL acetone 8 add 5.0 mL cyclohexane

4 add 7.0 mL acetone 9 add 4.0 mL cyclohexane

5 add 9.0 mL acetone 10 add 5.0 mL cyclohexane

Read this section carefully – a GA will demonstrate this portion of the Important: procedure in the laboratory, so be sure to be very familiar with the following.

To perform distillations, ensure that the top of the thermometer bulb is well below the sidearm leading to the condenser. Also, it is critical that the distillation proceeds at a slow rate so that both liquid and vapour are near thermal and compositional equilibrium. In order to do this, the Variac should be set between 60 and 90.

Note: For the pure cyclohexane sample, the Variac may be set near 90, but for other distillations, the Variac should be adjusted downwards.

It is important that cool water runs through the condenser throughout the experiment, and that the condenser is in its reflux position when heating is started for each mixture. Additionally, temperature must be monitored closely during each distillation.

Note: The flask can be “tapped” or “flicked” regularly to prevent superheating and bumping of solution. Please ask the GA to demonstrate this!

Detailed Procedure for each Distillation

1. When distillation is proceeding slowly (boiling has just begun, and the condenser is dripping), the condenser should be twisted into its collecting position.

2. When the condenser has collected a sufficient volume of distillate, the temperature should be recorded, and the heating mantle should be removed from below the flask (safety first!).

Experiment 5 –Liquid-Vapour Equilibrium in a Binary System

3. A cold water bath should then be used to replace the heating mantle. This will quickly cool the mixture and stop the distillation. This should be done by bringing the cold water bath to the boiling flask, and not vice versa, so that the distillate in the condenser is not spilled back into the mixture below.

4. Sample the liquid mixture in the boiling flask first. This is accomplished by inserting a syringe through the septum of the apparatus, and drawing up a small volume of the mixture.

5. The refractive index of this portion can be immediately determined using the Abbe refractometer (see below). This should be recorded as the “liquid” refractive index at the determined temperature.

6. Be sure to return the unused portion from the syringe back into the boiling flask (piercing the septum again).

Note: Ensure that the syringe is cleaned with acetone between samples.

7. The “vapour” (distillate) can then be sampled by inserting the syringe down into the condenser.

8. Measure the refractive index of the distillate. These refractive indices can be used, along with Table 5.2, to determine the vapour and liquid compositions in mole % for all ten mixtures.

9. Finally, return the unused distillate to the boiling flask, twist the condenser into its refluxing position, and proceed to the distillation of the next mixture. Discard all residues and distillates into the waste bottle when finished.

Refractometer Operation

1. Open the viewing chamber and clean it with a Kimwipe and acetone.

2. Spread a thin film of the solution on the glass plate with the syringe. Do not touch the plate with the syringe. Close the viewing chamber.

3. Rotate the large knob on the right until the dark/light separation is close to the diagonal crosshairs in the eyepiece.

4. Rotate the smaller knob on the right to fine tune the position of the light/dark separation.

5. Read the refractive index on the upper scale using the central vertical line as a reference point.

6. Clean the glass plate with a Kimwipe and acetone immediately after taking a reading.

Experiment 5 –Liquid-Vapour Equilibrium in a Binary System

Results

1. Determine the mole fraction for the liquid and vapour portions of each of the 10 mixtures by referring to Table 5.2.

2. On a single graph, make 2 plots of temperature as a function of the mole fraction of acetone. One of these plots will be for the “vapour” and the other will be for the “liquid”.

3. Determine the composition and boiling point of the azeotrope. As well, determine the boiling points of each of the pure liquids. Compare these values to those given in the literature.

4. Label the phases and components present in each area of the diagram, as well as the boiling temperatures of the pure compounds.

Lab Questions

1. Below is a table of common azeotropic mixtures, with temperatures given in °C:

* * System A B TA TB Taz (wt%B)az 1 1-propanol water 97.2 100 88.1 28.2 2 chloroform acetone 56.5 61.2 64.7 69.5 3 2-propanol cyclohexane 82.3 81.4 68.6 67

(a) Make a rough sketch of what you think the temperature-composition diagram might look like for each system. On your plot, the temperature will be on the y-axis and the composition (weight per cent of B) will be on the x axis. (b) Identify each system as a high-boiling or low-boiling azeotrope. (c) For system 3, at a composition of 75% cyclohexane, is it possible to distill pure cyclohexane? Explain why or why not.

References

1. Atkins, Peter. Physical Chemistry. 7th edition. Chapter 8 or 8th edition Chapter 6.

2. Shoemaker, David P. and Garland, Carl W. (1970). Experiments in Physical Chemistry. (McGraw-Hill).

3. Alberty, Silbey. (2001). Physical Chemistry. 3rd edition. Chapter 6, Section 6.5.

Recommended Reading

Partial pressures: Lecture 2; Raoult’s Law: Lecture 16

Lever rule: Lecture 18; Liquid-liquid phase diagrams and azeotropes: Lecture 19

Experiment 5 –Liquid-Vapour Equilibrium in a Binary System

Table 5.2

Mole% 25 Mole% 25 Mole% 25 Mole% 25 nd nd nd nd Acetone Acetone Acetone Acetone

0.00 1.4238 25.00 1.4091 50.00 1.3919 75.00 1.3744

1.00 1.4233 26.00 1.4085 51.00 1.3912 76.00 1.3737

2.00 1.4228 27.00 1.4079 52.00 1.3905 77.00 1.3730

3.00 1.4222 28.00 1.4072 53.00 1.3898 78.00 1.3723

4.00 1.4217 29.00 1.4066 54.00 1.3891 79.00 1.3715

5.00 1.4211 30.00 1.4059 55.00 1.3883 80.00 1.3708

6.00 1.4206 31.00 1.4053 56.00 1.3877 81.00 1.3700

7.00 1.4200 32.00 1.4047 57.00 1.3873 82.00 1.3693

8.00 1.4195 33.00 1.4042 58.00 1.3866 83.00 1.3685

9.00 1.4189 34.00 1.4035 59.00 1.3859 84.00 1.3678

10.00 1.4183 35.00 1.4029 60.00 1.3852 85.00 1.3671

11.00 1.4177 36.00 1.4021 61.00 1.3845 86.00 1.3664

12.00 1.4171 37.00 1.4015 62.00 1.3838 87.00 1.3657

13.00 1.4165 38.00 1.4008 63.00 1.3831 88.00 1.3649

14.00 1.4159 39.00 1.4001 64.00 1.3824 89.00 1.3642

15.00 1.4153 40.00 1.3994 65.00 1.3815 90.00 1.3635

16.00 1.4147 41.00 1.3989 66.00 1.3808 91.00 1.3628

17.00 1.4140 42.00 1.3981 67.00 1.3801 92.00 1.3621

18.00 1.4134 43.00 1.3974 68.00 1.3794 93.00 1.3614

19.00 1.4128 44.00 1.3966 69.00 1.3786 94.00 1.3607

20.00 1.4122 45.00 1.3959 70.00 1.3779 95.00 1.3600

21.00 1.4116 46.00 1.3951 71.00 1.3772 96.00 1.3593

22.00 1.4110 47.00 1.3943 72.00 1.3764 97.00 1.3587

23.00 1.4104 48.00 1.3936 73.00 1.3759 98.00 1.3579

24.00 1.4098 49.00 1.3926 74.00 1.3751 99.00 1.3572

100.00 1.3563