phy chem 2 lab

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ADAMSON UNIVERSITY

College of Engineering

Chemical Engineering Department

M a n i l a

Physical Chemistry for

Engineers 2

Laboratory Manual

Prepared by:

Committee for Laboratory Manual of the Chemical Engineering Department


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Table of Contents:

Course Description (i )

Policies and Guidelines on the use of ChE/Chem Laboratories ( ii )

Instruction for Preparing Laboratory Reports (i ii )

Experiment

No.

Title

1 Disti l lat ion of Binary Liquids

2 Steam Disti l lat ion

3 Solubil ity Diagram of a Part ial ly Miscible Liquid

System

4 Distr ibution

5 Phase Diagram of a Three Component Liquid System

6 Adsorption

7 Chemical Equil ibr ium

8 Chemical Kinetics

9 Galvanic Cells

10 Conductimetry


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COURSE SYLLABUS IN PHYSICAL CHEMISTRY FOR ENGINEERS -2

(Laboratory)

Course Description:

This laboratory course accompanying Physical Chemistry 2 (lecture) is a continuation of

Physical Chemistry 1 laboratory which covers the experiments on chemical equilibria, phase equilibria,

surface phenomena, thermochemistry, kinetics and electrochemistry.

No. of hours : 3 hours per week

Course Credit Units : 1 unit

Course Objectives:

After completing this course, the students must be able to:

1. Develop sound judgement in interpreting and correlating experimental data based on theprinciples learned in Physical Chemistry.

2. Develop initiative, resourcefulness, and leadership by demonstrating full responsibility inperforming the experiments assigned.

3. Acquire laboratory skills by following accepted laboratory handling and waste disposaltechniques.

4. Develop safety consciousness by observing proper laboratory techniques at all times whileworking in the laboratory.


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POLICIES AND GUIDELINES ON THE

USE OF CHEMICAL ENGINEERING/CHEMISTRY LABORATORIES

1. Chemical Engineering laboratories are open for equipment review and other preparatory work from 8

am to 5 pm Monday through Friday. Equipment may only be operated and chemicals will bedispensed during scheduled laboratory hours. Written approval of the faculty adviser, laboratory

coordinator, and ChE chairperson are required for anyone (e.g. research students) to use the laboratoryduring Saturdays or to operate equipment on non-scheduled laboratory days.

2. For the use of equipment, computers, and chemicals, students and faculty members must sign in the

log books for proper monitoring of the equipment.

3. No equipment is to be operated until the approval of the instructor and laboratory coordinator has beenobtained at the check-in meeting. Only the equipment pertaining to the assigned experiment is to be

operated. All members of a group are to be properly informed on the safety aspects of their assigned

experiment and to be familiar with the safety aspects of surrounding experiments. Before any

apparatus can be operated in the laboratory, the group must have a second on-site safety check-in.

4. Students are required to prepare handling and storage procedures of chemicals and materials to beused and waste disposal/treatment procedures approved by the faculty adviser and laboratory

coordinator prior to any experiment.

5. Transfer or movement of equipment or devices from the laboratories will not be allowed without

approval of advance notice (at least 2 days before the request schedule) from the laboratory

coordinator.

6. In borrowing glassware and accessories of the equipment, equipment and in using laboratory facility,students must accomplish Form A (Request for the Use of Laboratory Facilities and Equipment) with

the signature of the professor.

7. In case of breakages, damages, or losses:(A) A student accomplishes three copies of the Breakage Forms/Charge Slip Form (Form B) duty

signed by all members of the group.(B) The student requests the signature of the instructor and returns the form to the laboratory

personnel.(C) The laboratory personnel indicate the price of the damaged or lost item.

(D) The students are given one week to replace the items with the same brand or specifications. If the

students fail to replace the said item, the item is then forwarded to the Cashiers Office for

payment.(E) After payment, the students return the forms to the laboratory personnel with the photocopy of the

official receipt.(F) They must be advised to keep the original receipt for future reference.

8. Students should always wear their laboratory gowns or aprons when working in the laboratory. Safety

goggles and safety gloves must also be worn whenever applicable. Students must also observe properattire specifically closed footwear to protect them from spilled chemicals or hot fluids. Neckties,

dangling clothing or jewelry and other unsafe items are prohibited.


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9. Sleeping is prohibited in the laboratory. Violation of this rule will result in dismissal from the

laboratory for that day.

10. Horseplay of any sort is absolutely prohibited in the laboratory.

11. As a general rule, eating, drinking, chewing gum, and smoking are not allowed inside laboratories.

12. All safety precautions must be followed at all times.

13. No operating equipment will be left unattended. At least two members of the group must be present

while the equipment is operating.

14. The laboratory floor must be kept dry, clean, and uncluttered at all times. Any spills should be cleanedup immediately.

15. All injuries, accident, hazardous situation, losses, leaks, malfunctions or breakages must be reported to

the laboratory personnel or professor immediately.

16. All chemicals must be transported in a safety carrier. All mercury and alcohol thermometers and morethan one item of glassware must be transported in a bucket or other suitable container.

17. The students are expected to be familiar with the safety aspects of all chemicals used in the laboratory.

18. Listening to radios, walkman, MP3s, MP4s, etc. is strictly prohibited in the laboratory.

19. Playing computer games; using and recharging of cellular phones; or viewing DVDs is prohibited in

the laboratory.

20. Applying cosmetics are prohibited in the laboratory.

21. Precautions should be taken to prevent long hair from being entangled in moving parts of theequipment.

22. A violation notice will be issued by the Laboratory Coordinator or Laboratory Student Assistant or by

the assigned Laboratory Professor to any student found violating any of these rules and regulations.

23. Any SERIOUS VIOLATION of any of these safety rules or laboratory policies may lead to

immediate dismissal from the laboratory. A person who repeatedly disregards the safety rules or

laboratory policies for at least 3 times will be called in for disciplinary action with the ChEchairperson. A penalty that suits the violation may be imposed and, at the discretion of his/her

laboratory professor, the students grade may be severely affected.

Other policies may be given as situation arises and in consideration of our best interest. It is expected that

these sets of policies serve as a guide for us to work safely and efficiently.


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Instructions for Preparing Laboratory Reports

The report must be prepared and will be graded according to the following outline:

Subject Marks

1. Title page 2

2. Abstract 12

3. Introduction 3

4. Theoretical Background 5

5. Procedure 4

6. Results 8

7. Discussion of Results 24

8. Conclusions and Recommendations 8

9. Literature Cited 2

10 Nomenclature 2

Appendices

A1 Raw Data 5

A2 Analysis of data & Sample Calculation 20

Organization & neatness 5

Total: 100

1. Title Page

The Title page should be separate from the rest of the report. It should contain:

a. The name of the experiment

b. The number of coursec. The date when the experiment was run

d. The name of the writer and his co-workers, ID number, and his group number

e. The name of the instructor to whom the report is submittedf. The date of submission of report


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2. Abstract

The abstract should be informative, and should be written in about three to five sentences. It should

cover all phases of the investigation. It must include the following:a. An introductory statement about the subject matter

b. Briefly describe what was done.

c. Present some selected result (numerical values, if possible).

d. If possible, present some percentage errors in experimental results in comparison with theoreticalvalues. While writing the abstract, it should be kept in mind that you should not refer

to any graph or table.

3. Introduction

This section should include few sentences discussing the physical and/or chemical principles involved in

the experiment.

4. Theoretical Background

This section should include the theory behind the experiment. It should also contain all those equations,

which are used to acquire a certain result. Theoretical correlations, which are used for comparison withexperimental results, should also be included.

5. Procedure

Here, you should briefly describe the actual step-by-step procedure you followed in running the

experiment. It should be written in your own words, e.g. the needle valve was manipulated in order to

adjust the liquid flow rate.

6. Results

The results should be presented in the form or Tables or graphs. The Table should contain the resultsobtain from experiments and from theoretical knowledge. Comparisons should be

presented in terms, e.g. percent deviation.

7. Discussion of Results

In this section you should discuss you experimental results. Show how you make comparison with the

values obtained theoretically. Also discuss the deviation of experimental results from theoretical values.The possible source of errors should also be mentioned. If the results are obtained in terms of graphs,

then interpret them also.

8. Conclusions and Recommendations

Conclusions are the series of numbered sentences which answer the questions posed in the end of each

experiments. Conclusions should also include the errors between the experimental and theoreticalvalues. What you have learned from the experiment should be mentioned as well. Recommendations arethe proposals for future work, e.g. suggested changes in equipment, study of new variables, or possible

experiments in relative fields. Like the conclusions, the recommendations are usually listed by numbers,

and each consists of only a sentence or two.

9. Literature Cited


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Here, you should list the books, Journals articles, etc. used in writing your report and analyzing theexperiment. The reference should be completed (name of the book, author, volume, date of publication,

pages, etc.). References should be arranged alphabetically by author.

10. Nomenclature

The symbols, which are used in the report, should be defined in the nomenclature in alphabetical order.

The accompanying definitions must include proper units.

Appendices

All appendices and graphs should be attached at the end of the report.

A1. Raw Data:

It should contain the data on which the experiment was done.

A2. Analysis of Data and Sample Calculations:

This shows how the data/manipulated data is transformed into experimental results by using the

appropriate equations. Also, how the theoretical results are obtained using theoretical in terms ofpercentage error. Sample calculations should contain each step, which is used to acquire certain results.


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Experiment 1

DISTILLATION OF BINARY LIQUIDS

OBJECTIVE:

To construct a boiling point-composition diagram for a binary liquid system.

DISCUSSION:

When a mixture of two liquids is distilled, the boiling point will be that temperature at which the sumthe sum of the partial vapour pressures of the two liquids is equal to the atmospheric pressures. If the

liquids form an ideal solution, then the partial pressures can be calculated with the aid of Raoults law

and the total pressure can be expresses by the equation P=N1p1+ N2p2, where P is the total pressure, N1and N2 are the more fractions of compounds 1 and 2, respectively, and p1 and p2 are the vapour

pressures that the pure compounds 1 and 2, respectively, would have at the given temperature. If the

attractive forces between unlike molecules are greater than between like molecules, then the total

pressure will be less than that predicted by Raoults law; if the attractive forces between unlikemolecules are less than between like molecules, then the total pressure will be greater than that predicted

by Raoults law. A lower total pressure corresponds to a higher boiling point, and vice versa.

A boiling point composition diagram consists of two curves on the same graph. (1) a plot of the

composition (usually mole fraction) of a liquid against its boiling point, (2) a plot of the composition of

the vapour in equilibrium with the liquid at its boiling point.

MATERIALS:

250-ml distilling flask, condenser, adaptor, thermometer, 1225-ml Erlenmeyer flasks, 10-ml graduated

cylinder, 5-ml pipets, 100-ml beaker, 600-ml beaker, burette, iron stand, burette clamp, boiling chips,

glacial acetic acid, 1M NaOH, phenolphthalein indicator.

PROCEDURE:

1. Prepare a simple distillation set-up.

2. Mix 50-ml of glacial acetic acid and 1-ml of distilled water in a 100-ml beaker.

3. Pipet 1 ml of the mixture into each of two 125-ml Erlenmeyer flask and label both flasks L-1 to

make two trials.

4. Transfer the rest of the solution in step 2 to the distilling flask in the set-up and add the boiling

chips.

5. Slowly distill the mixture collecting the distillate in a clean, dry 10-ml graduated cylinder. Read

and record the temperature when 3-ml of the distillate has been collected. Stop the distillation

when 6-ml of the distillate has been collected.


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6. Pipet 1-ml of the solution remaining in the distilling flask into each of the Erlenmeyer flasks

previously labelled L-1 in step 3.

7. Pipet 2-ml of the distillate from step 5 into each of the two Erlenmeyer flasks and label both

flasks V-1.

8. Add 5-ml of distilled water to the remaining solution in the distilling flask. Distill over 1-2 ml to

wash out the condenser.

9. Stop heating and pipet 1-ml of the solution into each of two 125-ml Erlenmeyer flasks labelled

L-2.

10. Repeat steps 5 to 7 labelling the distillate fractions as V-2.

11. Repeat steps 8-10 using 10, 20, and 30 ml of distilled water labelling the flasks L-3 and V-3, L-4

and V-4, and L-5 and V-5, respectively.

12. Titrate the solutions in the flasks with 1 M NaOH using phenolphthalein as indicator.

13. Record the data on the report sheet and calculate the mole% of acetic acid in the solutions.

Assume that the density of all solutions is 1.05g/ml.

14. Construct the boiling point-composition diagram of the binary liquid system acetic acid-water by

plotting boiling point against composition in mole% acetic acid.


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Experiment No. 1


PRELAB EXERCISES

NAME: __________________________________________ DATE: _________________

PROFFESOR: ______________________________________ SCORE: ________________

1. At 90C, the vapour pressure of methylbenzene is 400 Torr and that of 1,2-dimethylbenzene is

150 Torr. What is the composition of a liquid mixture that boils at 90C when the pressure is

0.50 atm? What is the composition of the vapour produced?

2. The vapour pressure of pure liquid A at 300 K is 575 Torr and that of pure liquid is 390 Torr.

These two compounds form an ideal liquid and gaseous mixtures. Consider the equilibriumcomposition of a mixture in which the mole fraction of A in the vapour is 0.350. Calculate the

total pressure of the vapour and the composition of the liquid mixture.

3. It is found that the boiling point of a binary solution of A and B with xA=0.6589 is 88C. At this

temperature the vapor pressures of pure A and B are 957.0 Torr and 379.5 Torr, respectively. (a)

Is this solution ideal? (b) What is the initial composition of the vapour above the solution.


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Experiment No.1


PRELIMINARY DATA SHEET:

Date: _______________ Score: _______

Group No.: ___________

SIGNATURE

Leader: ____________________________________ __________________________

Members: ____________________________________ ____________________________

____________________________________ __________________________

_______________________________________ ____________________________

THEORETICAL FRAMEWORK: (PREPARE A FLOWCHART OF THE PROCEDURE OF THE

EXPERIMENT)

EXPERIMENTAL RESULTS AND OBSERVATIONS:

Run

No.

Trial

Boiling

Point

ml of NaOH Mole % HC2H3O2

1 2 1 1 1 2 1 2 Ave.

L V L V L V

1

2

3

4

5


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Concentration of NaOH solution _____________________ M

Boiling Point of HC2H3O2 _____________________

Boiling Point of H2O _____________________

Boiling Point Composition Diagram of the System HC2H3O2-H2O

Sample Computation:


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Approved by: _______________________ PROFESSOR

GUIDE QUESTIONS AND PROBLEMS:

1. How do attractive forces between like molecules compare with attractive forces between

unlike molecules in the system acid-water? Support your answer.

2. How do you expect the composition of the vapour above the solution of a mixture of acetic

acid and water to compare with that of the liquid with which it is in equilibrium?

3. Define an Azeotrope. Does the system form an azeotrope mixture? Explain your answer.


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Experiment No. 2

STEAM DISTILLATION

OBJECTIVE:

To study the process of steam distillation.

DISCUSSION:

When, in a mixture of two liquids, the attraction between like molecules is greater than that between

unlike molecules, the total vapour pressure is greater than that predicted by Raoults law. As this

difference in attraction increases, a point will be reached where the two liquids will no longer miscible

in all proportions; further increase will decrease the range of miscibility until, as a limiting case, the

liquids can be considered to be completely immiscible in each other. The total vapour pressure over

these systems will differ more and more from that predicted by Raoults law until, in the limiting case

mentioned above, the vapour pressure will be equal to the sum of the vapour pressures of the two

pure liquids at the same temperature and will be independent of the mole fraction of each liquid

present.

In this limiting case, the ratio of the partial pressures of the compounds in the vapour will be

equal to the ratio of the vapour pressures of the pure compounds, at a temperature equal to the

boiling point of the mixture. Thus,

= = = = =

and


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=

where pA and pB are the vapour pressures of the pure compounds, pA and pB are the partial pressures of

the compounds in the vapour, P is the total pressure, N A and NB are the mole fractions of the compounds

in the vapour (or distillate)m nA and nB are the moles of the compounds in the vapour (or distillate), w A

and wB are the weights of the compounds in the vapour (or distillate), and MA and MB are the molecular

weights of the compounds.

Steam distillation is found to be a useful method of distilling a liquid which is immiscible with water,

has an appreciable vapour pressure between 80 and 100C., and would decompose if distilled directly

from the mixture at the higher temperature required. Although toluene does not decompose, it fulfils the

other requirements and will be used here because of experimental convenience.

MATERIALS:

250-mL distilling flask, 1-L round-bottom flask, condenser, thermometer, glass tubing, 50-mL

graduated cylinder, adapter, toluene

1. Prepare a steam distillation set-up as in the figure below. Have the set-up approved by your

instructor before proceeding with the experiment.


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source: http://www.chembook.co.uk/chap23.htm

2. Place 50-mL of toluene in the distilling flask. Heat the steam generator using a large flame.When the water starts boiling lower the flame to a size just enough to continue boiling the water.

3. Discard the first 2-3 mL of distillate. Record the temperature before placing the receiver (50-mL

graduated cylinder) under the adapter.4. Continue the distillation and record the temperature when the total volume of distillate in the

receiver has reached 5mL, 10mL, 15mL, 20mL, 25mL, 30mL, 35mL, 40mL, 45mL and 50 mL.

Stop the distillation when 50mL of distillate has been collected.5. Cool the distillate and allow the layers to separate completely. Read and record the volumes of

the layers obtained.

6. From a handbook, secure data for the vapor pressure of toluene and water at 40C, 50C, 60C,

70C, 80C, 90C. Determine the vapor pressure of the mixture at given temperatures and plot the

vapor pressures of toluene, water and the mixture on the same graphing paper.7. From the graph obtained in step 6, determine the theoretical boiling point of the mixture.

8. Determine the theoretical composition of the distillate in mole % toluene and weight % tolueneusing the graph obtained from the previous step.

Experiment No. 2

STEAM DISTILLATON

PRELAB EXERCISES

NAME: __________________________________________ DATE: _________________

PROFFESOR: ______________________________________ SCORE: ________________

1. Benzene and toluene from nearly ideal solutions. At 20C the vapour pressures of pure benzeneand toluene are 74 Torr and 22 Torr, respectively. The solution is boiled by reducing the

external pressure below the vapour pressure. Calculate (a) the pressure when boiling begins, (b)

the composition of each component in the vapour, and (c) the vapour pressure when only a fewdrops of liquid remains. Assume that the rate of vaporization is low enough for the temperature

to remain constant at 20C.


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2. The following temperature/composition data were obtained for a mixture of two liquids A and B

at 1.00 atm, where x is the mole fraction in the liquid and y is the mole fraction in the vapour atequilibrium:

/C 125 130 135 140 145 150

xA 0.91 0.65 0.45 0.30 0.18 0.098

yA 0.99 0.91 0.77 0.61 0.45 0.25

The boiling points are 124C for A and 155C for B. Plot the temperature/composition diagram

for the mixture. What is the composition of the vapour in equilibrium with the liquid of the

composition (a) xA=0.50 and (b) xB=0.33?

Experiment No. 2

STEAM DISTILLATION

PRELIMINARY DATA SHEET

Date: _______________ Score: _______

Group No.: ___________

SIGNATURE

Leader: ____________________________________ __________________________

Members: ____________________________________ ___________________________

____________________________________ __________________________

____________________________________ __________________________


EXPERIMENT)

EXPERIMENTAL RESULTS AND OBSERVATION:


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Volume of distillate Boiling Point

0 mL _____

5 mL _____

10 mL _____

15 mL _____

20 mL _____

25 mL _____

30 mL _____

35 mL _____

40 mL _____

45 mL _____

50 mL _____

Temperature ( ) 40 50 60 70 80 90

Vapor pressure of water (Torr) _____ _____ ______ ______ ______ _____

Vapor pressure of toluene (Torr) _____ _____ ______ ______ ______ _____

Vapor pressure of the mixture (Torr) _____ _____ ______ ______ ______ _____

Plot of P vs. T for Toluene, Water and the Mixture


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Experimental boiling point of the mixture ________

Theoretical boiling point of the mixture ________

% Error ________

Volume of toluene in distillate _______mL

Volume of water in distillate _______mL

Temperature of distillate _________

Density of water (from handbook) ______g/mL

Density of toluene (from handbook) ______g/mL

Weight of water in distillate ________g

Weight of toluene in distillate ________g

Moles of water in distillate _______moles

Moles of toluene in distillate _______moles

Experimental composition of distillate _____mole % water


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_____mole % toluene

_____wt % water

_____wt % toluene

Theoretical composition of distillate _____mole % water

_____mole % toluene

_____wt % water

_____wt % toluene

% Error _____%

SAMPLE COMPUTATIONS:

Approved by: _______________________

PROFFESOR


1. Give the advantages of using steam distillation in the recovery of an inorganic liquid from a

mixture.


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2. Is it possible to carry out steam distillation at a temperature higher than 100C at 1 atm? Why?

Experiment No. 3

SOLUBILITY DIAGRAM

OF A

PARTIALLY MISCIBLE LIQUID SYSTEM

OBJECTIVE:

To be able to prepare the solubility diagram of a partially miscible liquid system.

DISCUSSION:

In most cases, partially miscible liquids become more soluble in each other with increasing temperature

until the critical solution temperature is reached, above which there is complete miscibility. In thisexperiment mixtures of known composition are heated until they dissolve completely. This is easily

recognized by the disappearance of cloudiness due to two phases. A plot of composition vs. temperature

shows the critical solution temperature.

The phenol-water system readily gives satisfactory data. Since pure phenol is solid at room temperature

and is messy to handle, an 80 percent by weight solution is used. This is dispensed from a burette, and


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the weight is calculated from the volume and the density. Any solution of phenol is dangerous, so be

careful. Avoid spilling and clean up immediately with large amounts of water. Never pipette by mouth.

MATERIALS:

50-mL test tube, acid buret, iron stand, buret clamp, copper wire stirrer, thermometer, 600-mL beaker,

tripod, burner, melted phenol

PROCEDURE:

1. Prepare 80% phenol solution as follows:

a. Weigh a clean and dry 100-mL graduated cylinder.

b. Place about 20-mL of melted phenol in the cylinder and weigh. Record the weight of thephenol.

c. By means of a buret, add distilled water equivalent to the weight of the phenol in the

cylinder.

(CAUTION: Phenol is corrosive. Avoid contact with your skin.)

2. Transfer all the phenol solution to a clean acid buret.

3. Deliver 10 mL of the phenol in the buret into a 50-mL test tube. Cover with a cork fitted with a

thermometer and a copper wire stirrer.

4. By means of another buret, add distilled water to the phenol solution stirring continuously ring

addition until the solution turns cloudy. Record the volume of the phenol solution.

5. Immerse the tube in a water bath (use a 600-mL beaker), then heat the water in the bath until the

solution turns clear. Note the clearing temperature. (NOTE: The solution in the tube must becontinuously stirred while doing this step.)

6. Remove the tube from the water bath and continuously stir the solution until the solution turns

cloudy. Record the clouding temperature.

7. Deliver an additional 1.0 mL of water from the buret then repeat steps 5 and 6.

8. Repeat step 7 four (4) times.

9. Deliver 10.0 mL of distilled water from a buret into another ignition tube.


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10. To the distilled water in the tube, add phenol solution from a buret while stirring constantly untilthe solution turns cloudy.

11. Treat the solution in the tube as in steps 5-8.

12. Record all data in the table.

13. Calculate the % by weight of pure phenol in all solutions obtained. The density of the 80%

phenol solution is 1.05 g/mL.

Experiment No. 3

SOLUBILITY DIAGRAM

OF A


PRELAB EXERCISES

NAME: __________________________________________ DATE: _________________

PROFFESOR: ______________________________________ SCORE: ________________

The figure below shows the phase diagram for two partially miscible liquids, which can be taken to be

that for water (A) and 2-methyl-1-propanol (B). Describe what will be observed when a mixture of


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composition x_b=0.8 is heated, at each stage giving the number, composition, and relative amounts ofthe phases present.

0 0.2 0.4 0.6 0.8 1

Experiment No. 3

SOLUBILITY DIAGRAM

OF A



Date: _______________ Score: _______

Group No.: __________

SIGNATURE

Leader: ____________________________________ __________________________

T1

Temperature,

T


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Members: ____________________________________ __________________________

____________________________________ __________________________

____________________________________ __________________________


EXPERIMENT)


A. Water Add to Phenol

Run No. 1 2 3 4 5 6

Clearing temperature (oC)

Clouding temperature (oC)

Average temperature (oC)

Volume of phenol solution

Weight of the phenol solution (W1)

Weight of pure phenol

Weight of water added to phenol (W2

)

Total weight of solution (WT)

Weight % phenol

B. Phenol Added to Water

Run No. 1 2 3 4 5 6

Clearing temperature (oC)

Clouding temperature (oC)

Average temperature (oC)

Volume of phenol solution

Weight of phenol solution (W1)


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Weight of pure phenol

Weight of water added to phenol (W2)

Total weight of solution (WT)

Weight % phenol

COMPUTATIONS FOR A AND B:

Experiment No. 4

DISTRIBUTION

OBJECTIVE:

To determine the values of K and n in the distribution of acetic acid between water and diethyl ether.

DISCUSSION:

When a solute is added to a mixture of two miscible liquids and the system allowed to come to

equilibrium at some given temperature, the following relationship will hold regardless of the quantity of

solute added:


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Ka=

where Ka is the distribution constant, a1 and a2 are the activities of the solute in the two solvents, and n isthe ratio of the average molecular weight in solvent 1 to its average molecular weight in solvent 2.

Association will increase the average molecular weight, while dissociation will decrease it. It is to benoted that n may vary considerably with concentration due to changes in the degrees of association anddissociation.

In order to simplify this experiment, concentrations will be substituted for activities in the above

equation, so that the distribution equation becomes

Kc=

In this equation Kc may vary appreciably due to the change in activity coefficients with concentration.

If, for any distribution system, the values of C1 and C2 are obtained for two different total soluteconcentrations, both n and K can be calculated by solving the two equations for the two unknowns. In

carrying out these calculations it may be convenient to take the logarithm of both sides of the above

equation. The values of K and n obtained will be the average values for the two different concentrationsin the calculations.

MATERIALS:

Glacial acetic acid, NaOH pellets, distilled water, phenolphthalein, Erlenmeyer flasks, base burette, iron

stand, burette clamp, separatory funnel, iron ring, pipettes.

PROCEDURE:

1. Prepare by serial dilution 60 mL of each of the following starting from glacial acetic acid:1.0 M, 0.50 M, and 0.25 M. (NOTE: to prepare solutions by serial distillation, prepare 120

mL of the 1.0 M solution from glacial acetic acid which is 17 M. Next, prepare 120mL of thesolution from the 1.0 M solution. Finally, prepare 120 mL of the 0.25 M solution from 0.50M solution.) Calculate the volumes of the acetic acid solution and water needed to prepare

the desired solution. )

2. Prepare from NaOH pellets 0.50 M NaOH solution then prepare by distillation 200mL of

0.10 M NaOH solution.


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3. Pipet 25mL of the 1.0 M solution into a separatory funnel. Add 25mL of diethyl ether shakewell and let stand for 10 minutes or until the layers are completely separated.

4. Draw the aqueous layer into a clean container then pipet 5.0mL into a clean 125-mL

Erlenmeyer flask. Add 10mL of distilled water then titrate with 0.50 M NaOH to the

phenolphthalein endpoint. Make 2 trials.

5. Pipet 10mL of the ether layer into 125-mL Erlenmeyer flask and add 10mL of distilled water.

Titrate with 0.10 M NaOH to the phenolphthalein end point. Make 2 trials.

6. Repeat steps 3 to 5 using the 0.50 M and 0.25M acetic acid solution. Pipet 10mL of the

aqueous solution for titration.

7. Calculate the concentration of the acetic acid in each solution.

8. Determine the values of K and n for the concentration used by the equation:

K=

Where; = concentration of the solute in the aqueous solution

= concentration of the solute in the organic solvent

n = ration of the average molecular weight of the

solute in the solvent 1 to that in solvent 2If the solute dissociates in aqueous solution, its average molecular weightin water decreases and n is greater than 1.

If the solute associates in water solution, its average molecular weight in

water increases and n is less than 1.

9. Record the results in the table.

Experiment No. 4

DISTRIBUTION

PRELAB EXERCISES:

NAME: __________________________________________ DATE: _________________

PROFFESOR: ______________________________________ SCORE: ________________

1. Discuss Nernst Distribution Law, its application and limitations.


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Experiment No. 4

DISTRIBUTION


Date: _______________ Score: _______

Group No.: ___________

SIGNATURE

Leader: ____________________________________ __________________________


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Members: ____________________________________ ____________________________

____________________________________ __________________________

____________________________________ __________________________


EXPERIMENT)


A. Initial Concentration of Acetic Acid: 1.0M

Solvent Trial V NaOH MNaOH Vsolution Msolution

Water

1 ___________ ___________ __________

_

__________

_

2___________ ___________ __________

_

__________

_

Average xxx xxx xxx __________

_

Ether

1___________ ___________ __________

_

__________

_

2___________ ___________ __________

_

__________

_

Average Xxx xxx xxx __________ _

B. Initial Concentration of Acetic Acid: 0.5 M


1___________ ___________ __________ __________


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Water_ _

2

___________ ___________ __________

_

__________

_


_

Ether

1___________ ___________ __________

_

__________

_

2___________ ___________ __________

_

__________

_


_

C. Initial Concentration of Acetic Acid: 0.25 M


Water

1

___________ ___________ ___________

___________

2

___________ ___________ ___________

___________

Average Xxx xxx xxx __________

_

Ether

1

___________ ___________ __________

_

__________

_2

___________ ___________ __________

_

__________

_

Average Xxx xxx xxx __________

_


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COMPUTATIONS FOR A, B AND C:


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Concentration Range K n

0.50 1.0 M _____________ _____________

0.25 0.50 M _____________ _____________

GUIDED QUESTIONS AND PROBLEMS:

1. Does aetic acid dissociate or associate in water? Why?

2. How does dilution affect dissociation/association of acetic acid in water? Explain your answer.


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Experiment 5

PHASE DIAGRAM OF A

THREE-COMPONENT LIQUID SYSTEM

OBJECTIVE:

To prepare the phase diagram of a three-component liquid system.

DISCUSSION:

The composition of a ternary system may be described by one point in a triple coordinate

diagram. The phase diagram of a ternary liquid system separating into two phases is given in

Figure 1. The points on the dome (curve abcdefg), represent the compositions at which the two

phases separate.

C

0.001.0

0.25

0.8

0.50

0.6

d

c

0.75

b

0.4

e

f0.2

1.00 ag 0.0

A 0.0 0.2 0.4 0.6 0.8 1.0 B

Figure 1. Phase diagram of a ternary system with two immiscible liquids, A and B.


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Above this dome only a single phase, hence complete miscibility exists. At a composition

described by a point under this dome, the system will separate into two phases. The a and g

positions in Figure 1 indicate that there is slight miscibility between components A and B. If

no miscibility existed between components A and B, the a position would coincide with

corner A and the g position with corner B. The diagram also indicates that it is the third

component, C, that is really miscible with either A or B in all proportions. In such a phase

diagram the tie lines have a very important aspect: they connect the concentration of the two

phases experimentally found to be in equilibrium with each other. For instance, when a

mixture with composition h (Fig. 1) is prepared, it separates into two phases. Phase one (rich

in A and C, and poor in B) has the composition designated on the diagram by b. Phase two

(rich in B and C, but poor in A) has a composition designated by point f. The quantitative

ratio of the two phases is given by,

Phase_one fh

Phase_two bh

and, therefore, once a phase diagram is available it can be used to determine the compositions

and proportions of the phases that would result when a mixture of specified overall

composition is prepared.

You may notice that the dome in Figure 1 is not symmetrical and the tie lines are not

parallel to each other. This is simply because the solubility of C in the two phases (A and B)

is not the same. In whatever direction the tie lines are slanted, they connect points of

equilibrium compositions. These equilibrium compositions, b vs. f and c vs. e, become

increasingly similar with each subsequent tie line, starting from the base of the dome and

proceeding upward. Similarly, the tie lines become shorter and finally converge to a

composition. This is called isothermal critical point or the plait point.


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MATERIALS:

CHCl, Glacial acetic acid, NAOH pellets, acid burets, 125-ml Erlenmeyer flasks, pipets

PROCEDURE:

A. Preparation of Solubility Curve:

1. Prepare 10-ml mixtures of H2O and CHCl containing the following % by volume of

water in 125-ml Erlenmeyer flasks: 10%, 20%, 30%, 40%, 50%, 60%, 70%, 80% and

90%. Measure the volume of each liquid accurately by using burets.

2. Titrate each solution from step 1 above with glacial acetic acid until the solution no

longer turns cloudy when shaken vigorously.

3. Record the volume of each liquid in the table.

4. Calculate the weight of each liquid in each solution.

5. Calculate the % by weight of each liquid in each solution.

6. Plot the composition of the mixture on a Stokes and Roozeboom diagram (equilateral

triangle) and connect the points to construct a solubility curve. Extrapolate the curve

to zero on both ends.

B. Construction of Tie-Lines

1. Prepare 20-ml mixtures of the three liquids with the following composition (% by

volume):

Mixture 1: 15% HOAc 25% CHCl 60% HOMixture 2: 30% HOAc 25% CHCl 45% HO

Mixture 1: 45% HOAc 25% CHCl 30% HO

Mixture 1: 60% HOAc 25% CHCl 15% HO

If no layers are formed, increase the amount of water until a considerable amount of the

second layer can be observed.


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2. Transfer each mixture into a separatory funnel, shake well for 2 minutes and allow the

layers to separate.

3. Pipette 5.0 mL of each aqueous layer into previously weighted 250-mL Erlenmeyer

flasks, then weigh again, Add 50 mL of distilled water and titrate to the phenolphthalein

end point with 0.5 M NaOH.

4. Pipette 5.0 mL of each CHCl3 layer into previously weighted 250-mL Erlenmeyer flasks

and weigh again. Add 50 mL water and titrate to the phenolphthalein end point with 0.25

M NaOH.

5. Calculate the percent by weight of acetic acid in each layer.

6. Locate the point on the solubility curve and construct the tie lines by connecting thecompositions of the complementary layers with a straight line.

7. Determine the plaint point of the system.


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Experiment No. 5

PHASE DIAGRAM OF A

THREE-COMPONENTLIQUID SYSTEM

PRELAB EXERCISES:

NAME: __________________________________________ DATE: _________________

PROFFESOR: ______________________________________ SCORE: ________________

1). Define the following terms:

a). Plait Point

b). Solubility Curve

c). Ternary Phase Diagram (Stokes and Roozeboom)d). Tie Lines


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Experiment 5

PHASE DIAGRAM OF A

THREE-COMPONENTLIQUID SYSTEM


Date: _______________ Score: _______

Group No.: __________

SIGNATURE

Leader: ____________________________________ __________________________

Members: ____________________________________ __________________________

____________________________________ __________________________

____________________________________ __________________________

THEORETICAL FRAMEWORK: (PREPARE A FLOWCHART OF THE PROCEDURE

OF THE EXPERIMENT)


Mixture Volume of liquid Weight of liquid % by Weight

dD

HO HOAc CHCl

_____ _____ _____

_____ _____ _____

_____ _____ _____

_____ _____ __________ _____ _____

_____ _____ _____

_____ _____ _____

_____ _____ _____

_____ _____ _____

HO HOAc CHCl

_____ _____ _____

_____ _____ _____

_____ _____ _____

_____ _____ _____

_____ _____ _____

_____ _____ _____

_____ _____ _____

_____ _____ _____

_____ _____ _____

No.

1

2

3

4

5

6

7

8

9

HO HOAc CHCl

_____ _____ _____

_____ _____ _____

_____ _____ _____

_____ _____ __________ _____ _____

_____ _____ _____

_____ _____ _____

_____ _____ _____

_____ _____ _____


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Density of HO:__________Density of HOAc:____________Density of CHCl:___________

B. Tie Lines:

Mixture No. 1 2 3 4

Aqueous Layer Weight __________ __________ __________ __________ V of NaOH __________ __________ __________ __________

% HOAc __________ __________ __________ __________

CHCl3 Layer Weight __________ __________ __________ __________

V of NaOH __________ __________ __________ __________

% HOAc __________ __________ __________ __________

Draw a Ternary Phase Diagram of Water-Acetic Acid-Chloroform System: (Separate Sheet

of Paper)


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Experiment 6

ADSORPTION

OBJECTIVES:

1) To determine the saturation value for monomolecular coverage for the adsoprtion of

acetic acid by activated charcoal

2) To determine the values of the constant k and n in the Freundlich equation.3) To determine the values of the constant a and b in Langmuirs equation.

DISCUSSION:

Adsorption is a surface phenomenon. It consists of an interaction between the moving molecules

of a liquid or gas phase and the relatively fixed molecules comprising a surface or interface. (Inthis discussion the term molecule will be used generically to include atoms and ions as well as

molecules). The moving molecules become bound to the surface more or less strongly.

Adsorption processes may be classified as physical or chemical. Another classification is

reversible and irreversible, referring to the comparative ease of removal, desorption of thebound molecules. Generally, physical adsorption (physisorption) is reversible, while the

chemical type (chemisorptions) is irreversible. The adsorbed molecules may come from a gas

phase, a liquid or from a solution in a liquid. The adsorbing surface may be either liquid or solid.

Adsorption from solution is usually monomolecular (monolayer coverage); i.e. adsorption ceases

when the surface is completely covered with a layer of one molecular thick. Adsorbed layersmore than one molecular thick have been proved to exist in certain cases as yet rare. The amount

of adsorption varies with the concentration of the solution. There are two equations for this

relationship for monomolecular adsorption. One is the empirical Freundlich equation:

= KCn or log ( ) = log k + n log C

in which x/m is the weight of the absorbed material per gram of adsorbent and C is the

concentration of the solution at equilibrium. The term k and n are constants to be evaluated fromexperimental data in each case.

in which x/m and C have the same significance as above, a is a constant proportional to the heatof adsorption and the temperature, and b is the amount of adsorption when the surface is covered

with a monomolecular film. Langmuirs equation is based on the assumption that the solid

surface is completely uniform. Deviations from the equation indicate lack of uniformity in the

surface.


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PROCEDURE:

1) Prepare 100 mL each of the following solutions by serial starting from glacial aceticacid(17 M): 1.0 M, 0.50 M, and 0.0625 M.

2) Measure 50 mL of the 1.0 M solution into a 125 mL Erlenmeyer flask, add 1.0 gram

of accurately weighed activated charcoal, shake the mixture and allow to equilibriatefor one hour. Agitate the mixture from time during the equilibriation process.

3) While equililbriation is taking place, pipet 5.0 mL of the original solution into a 250

mL Erlenmeyer flask and titrate with standard 1.0 M NaOH to the phenolpthalein end

point. Calculate the original concentration of the solution from the titration data.4) After equilibriation, filter off the activated carbon and titrate 5.0 mL of the filtrate

with 0.1 M NaOH to the phenolphthalein end point.

5) Repeat steps 2 to 4 using the other solutions prepared in step 1. For the titration, use

10.0 mL for the last three soultions.6) Calculate the equilibrium concentration (C2) and the specific adsorption (y) for each

concentration.

7) Plot y versus C1 and determine the saturation value for monomolecular coverage.

8) Plot log y versus log C1 to determine the saturation value for monomolecular

coverage.9) Plot (C1/y) versus C1 to determine a and b in Langmuirs Equation.


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Experiment No. 6

ADSORPTION

PRELAB EXERCISES:

NAME: __________________________________________ DATE: _________________PROFFESOR: ______________________________________

1). Distinguish between the following adsorption isotherms: Langmuir, BET, Temkin and

Fruendlich.

2). The data given below for the adsorption of CO on charcoal at 273 K. Confirm that they fitthe Langmiur isotherm, and find the constant K and the volume corresponding to complete

coverage. In each case V has been corrected to 1.00 atm (101.325 kPa).

P(kPa) 13.3 26.7 40.0 53.3 66.7 80.0 93.3

V(cm3) 10.3 19.3 27.3 34.1 40.0 458.5 48.0


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Experiment 6

ADSORPTION


Date: _______________ Score: _______

Group No.: ___________

SIGNATURE

Leader: ____________________________________ __________________________

Members: ____________________________________ __________________________

____________________________________ __________________________

____________________________________ __________________________


OF THE EXPERIMENT)


Initial Conc (C1) Final Conc (C2) Wacetic acid mcharcoal y

____________ ____________ ____________ ____________ ____________

____________ ____________ ____________ ____________ ____________

____________ ____________ ____________ ____________ ________________________ ____________ ____________ ____________ ____________

____________ ____________ ____________ ____________ ____________

SAMPLE COMPUTATION:


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Plot y versus C1

Saturation value for y monomolecular coverage = _____________________________


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Plot of log y versus C1

Freundlichs constant:

n =___________________________

k=____________________________


48/88

Plot of log (C1/y) vs C1

Langmuirs constants:

a = ______________________________

b =______________________________


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Experiment 7

CHEMICAL EQUILIBRIUM

OBJECTIVE:

To be able to determine the equilibrium constant for the system:

Fe3+ + CNS F eCNS2+

Yellow colorless blood-red

DISCUSSION:

The word equilibrium suggests balance or stability. The fact that a chemical reaction occurs

means that the system is not in equilibrium. The process will continue until the system reachesequilibrium. At this point, there is no observable change in the concentrations of reactants and

products. While the reaction appears to stop, the reality is that reactants are being converted to

products at the same rate that products are being converted toreactants.

These principles can be illustrated with a generic reaction:

a A + b B c C + dD

If only A and B are present, the reaction will proceed to the right; if only C and D are present,

the reaction will proceed to the left. In both cases, the reaction will apparently stopwhen equilibrium is reached.This state is characterized by a specific value for the

equilibrium constantK, defined for this reaction by the following expression.

K=

Here the brackets signify the equilibrium concentrations (in M) of the various species. Note that

the coefficients in the balanced equation appear as exponents in the equilibrium constant and thatthe products always appear in the numerator and the reactants in the denominator.

The value of the equilibrium constant depends on the chemical reaction and on temperature.

However, the value ofKwill n o t depend on the initial concentrations of reactants and products.

Furthermore, the specific mathematical form of the equilibrium constant must correspond to thecorrect chemical equation for the reaction. If the equation is not correct for example, if the


50/88

formula of one or more of the species is wrong the value ofK, even if calculated using accurate

concentrations and making no numerical errors, will not be constant.

In this experiment you will investigate the reaction of the Fe3+

[iron(III), or ferric] ion with

the SCN

(thiocyanate) ion. The product of this reaction is a complex ion that imparts a redcolor to aqueous solution. You are asked to experimentally determine the formula for the

complex ion and a numerical value for the equilibrium constant for the reaction in which it isformed.

PROCEDURE:

1. Label 5 clean and dry test tube of identical size and make.

2. Prepare solutions of Fe(NO3)3 by serial dilution as follows:

a.) Place 20 mL of 0.20 M Fe(NO3)3 in a 25-mL graduated cylinder.

b.) Transfer 8 mL of the solution into test tube 1.

c.) Add distilled water to the remaining solution in the graduated cylinder to bring thevolume to 20 mL.

d.) Transfer 8 mL of diluted solution into test tube 2.

e.) Repeat procedures c and d until 5 different concentrations of Fe(NO3)3 are obtained.

3. Add 8 mL of 0.002 M KCNS to each diluted solutions from step 2.

4. Wrap each test tube with white paper.

5. By looking down into the solutions in the test tube, compare the color intensities of thesolutions in test tubes 1 and 2. (The process may be aided by placing a mirror below the

test tubes).

6. Adjust the height of the solution in test tube 1 by removing or adding some solution untilthe observed color intensities appear to be equal. Transfer the solution in test tube 1 to

clean container with the aid of dropper.

7. Measure the heights of the solutions in test tubes 1 and 2 by means of a ruler.


51/88

8. Repeat steps 5, 6 and 7 with the test tubes 3, 4 and 5.

9. Record all measured heights of solution.

10. By assuming that all the limiting reagent in test tubes 1 was converted to the complex,calculate the concentration of the complex in test tubes 2 to 6 by using test tube 1 as

standard and by applying Beer-Lamberts law:

CS x hs = CU x hu

Where: CS=concentration of the standard

hs= height of the standard

CU=concentration of the unknown

hU=height of unknown

11. Calculate the equilibrium concentrations of the reactant.

[Reactants] equilibrium = [Reactants]initial [Reactants]converted

12. Record all the data in the table.

Experiment No. 7


PRELAB EXERCISES:

NAME: __________________________________________ DATE: _________________

PROFFESOR: ______________________________________

1. Discuss Beer-Lamberts Law, its application and limitations.


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Experiment 7



Date: _______________ Score: _______

Group No.: ___________

SIGNATURE

Leader: ____________________________________ __________________________


53/88

Members: ____________________________________ __________________________

____________________________________ __________________________

____________________________________ __________________________


OF THE EXPERIMENT)


Test Tube No. 1 2 3 4 5

Before [Fe3+]

________ ________ ________ ________ ________

Mixing [CNS-]________ ________ ________ ________ ________

After [Fe3+

]________ ________ ________ ________ ________

Mixing [CNS-]

________ ________ ________ ________ ________

hs (height of standard)xxx ________ ________ ________ ________

hu (height of unknown)

xxx ________ ________ ________ ________

[Fe3+]xxx ________ ________ ________ ________

[CNS-]

xxx ________ ________ ________ ________ [FeCNS2+]

________ ________ ________ ________ ________

keqxxx ________ ________ ________ ________

Computation of Average Keq:


1. How is the value of the equilibrium constant affected by the concentration of the

reactants used?


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2. How do you explain the above observation?

3. What factors affect the value of the equilibrium constant?

Experiment 8

CHEMICAL KINETICS

OBJECTIVES:

1. To determine the order of the reaction

2. To determine the activation energy of both uncatalyzed and catalyzed reactions.

DISCUSSION:

Chemical reactions occur at varying speeds with a vast spectrum of rates, ranging from very

slow to extremely fast. For example, the rusting of iron is fairly slow, whereas the decomposition


55/88

of TNT proceeds explosively fast. Experiments have shown that the rate of a homogeneous

reaction in solution depends upon the nature of the reactants, their concentrations, the

temperature of the system, and the use of catalysts.

Consider the hypothetical reaction:

A + B C + D

The rate of this reaction may be measured by observing the rate of disappearance of the reactantsA or B, or the rate of appearance of the products C or D. Which species is observed is a matter

of convenience. For example if A, B, and C are colorless and D is colored, the rate of appearance

of D can be conveniently measured by observing an increase in the intensity of the color of the

solution as a function of time. Mathematically, the rate of reaction may be expressed as follows

Rate of disappearance of A = = -

Rate of appearance of D = =

In general, the rate of the reaction depends upon the concentration of one or more of the

reactants. Thus, the rate of the hypothetical reaction above may be expressed as:

Rate = k[A]x

[B]y

where [A] and [B] are the molar concentrations of A and B, x and y are the powers to whichthe respective concentrations must be raised to describe the rate, and k is the specific rate

constant. The values of x and y must be determined experimentally. For example, if x = 2 and y =

1, then the rate law is:

Rate = k[A]2

[B]

The rate of reaction depends on concentration, temperatures, catalysts and inhibitors.

MATERIALS:

Zn metal, concentrated HCl, 1M Cu(NO3)2 solution, 30-mL test tubes, stopper fitted with right-

angle delivery tube, water bath.

PROCEDURE:

A. Effect of Concentration


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1. Prepare by serial dilution from concentrated HCl, 5M, 4M, 3M, 2M and 1M solutions.

2. Place 10-mL of 1M HCl in a clean and dry 30-mL ignition tube, add a piece of Zn metaland close with a stopper fitted with a right-angle and delivery tube.

3. Dip the tip of the delivery tube in a beaker of water and measure the time required for 30

bubbles to emerge from the tip of the delivery tube.4. Change the ignition tube with a clean and dry one, place 10-mL of 2M HCl and add Zn.

(You may use the unreacted Zn from the previous step after rinsing it well.). Repeat steps

2 and 3.5. Do steps 2 and 3 using 3M, 4M and 5M HCl.

6. Formulate the rate law. Determine the values of k and n in the rate law.

B. Effect of Temperature

1. Place 10-mL of 2M HCl in a clean, dry ignition tube and warm in a water bath until the

temperature of the HCl solution reaches 40C. Remove the tube from the water bath, adda piece of Zn and fit the tube with the delivery tube.

2. Determine the time required for 30 bubbles to emerge from the tip of the delivery tube.3. Repeat steps1 and 22 using temperatures of 50C, 60C and 70C.

4. Determine the values of k from the rate using data obtained from procedure A.

5. Plot log k versus 1/T and determine the activation energy of the reaction by using theArrhenius equation:

C. Effect of Catalyst

1. Place 10-mL of 2M HCl and 2-mL of 1M Cu(NO3)2 in a clean and dry ignition tube. Heat

the solution in the tube to 40C in a water bath. Remove the test tube from the water bath,add a piece of Zn and fit with the delivery tube.

2. Determine the time required for 30 bubbles to emerge from the tip of the delivery tube.

3. Repeat steps 1 and 2 at 50C, 60C and 70C.4. Determine the values of k from the rate using data obtained from procedure A.

5. Plot log k versus 1/T and determine Ea for the catalyzed reaction.

Experiment No. 8

CHEMICAL KINETICS

PRELAB EXERCISES:

NAME: __________________________________________ DATE: _________________


57/88

PROFFESOR: ______________________________________

1. The rate constant for the first order decomposition of a compound A in the reaction 2AP is k=2.78 x10-7 s-1 at 25C. What is the half life of A? What will be the pressure,initially 32.1 kPa, (a) 10 s, (b) 10 minutes after initiation of the reaction?

2. The rate constants for the decomposition of certain substance is 1.70 x 10-2 L mol-1s-1 at24C and 2.01 x 10-2 L mol-1s-1 at 37C. Evaluate the Arrhenius parameters of the

reaction.


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Experiment 8

CHEMICAL KINETICS


Date: _______________ Score: _______

Group No.: ___________

SIGNATURE

Leader: ____________________________________ __________________________

Members: ____________________________________ __________________________

____________________________________ __________________________

____________________________________ __________________________


OF THE EXPERIMENT)


A. Effect of Concentration

Concentration time Rate (1/time)

1M ______________ ____________

2M ______________ ____________

3M ______________ ____________

4M ______________ ____________

5M ______________ ____________


59/88

Plot of Rate vs. [HCl]

Rate law:

k = n=

B. Effect of Temperature

t (C) T(K) time k

40 _____ _____ _____ _____ _____

50 _____ _____ _____ _____ _____

60 _____ _____ _____ _____ _____

70 _____ _____ _____ _____ _____


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Plot of log k vs. 1/T

Rate law:

slope = Ea=

C. Effect of Catalyst

t (C) T(K) time k

40 _____ _____ _____ _____ _____

50 _____ _____ _____ _____ _____

60 _____ _____ _____ _____ _____

70 _____ _____ _____ _____ _____


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Plot of log k vs. 1/T

Rate law:

slope = Ea=

GUIDED QUESTIONS AND PROBLEMS:

1. Why is the reaction rate independent of the amount of Zn used in the experiment?

2. How does trebling the concentration of HCl affect the reaction rate?


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3. How does the activation energy of the catalyzed reaction compare with that of the

uncatalyzed reaction?

Experiment 9

GALVANIC CELLS

OBJECTIVES:

1. To be able to assemble galvanic cells.

2. To be able to measure the potentials to test electrodes.

3. To be able to measure the potential of a galvanic cells.

DISCUSSION*:

Oxidation and reduction reaction occur simultaneously side by side. A reduction reaction occurs

only if an oxidation reaction occurs and vise versa. Electrons are given in oxidation while in

reduction electrons are gained. Oxidizing agent is a chemical substance which has a large

tendency to gain electrons, while reducing agent is a chemical substance causes other substances

to be reduced and itself oxidized.

Example for an oxidation-reduction reaction:

2Ag+(aq) + Cu(s) 2Ag(s) + Cu2+(aq)

We can see that Cu(s) converts to Cu2+ (Oxidation). While Ag+(aq) changes to Ag(s) (reduction).

The oxidation-reduction reaction can be divided into two half reactions, one for oxidation and

the other is reduction reaction:

Cu(s) Cu2+(aq) + 2e- (Oxidation)

2Ag+(aq)+ 2e- 2Ag(s) (Reduction)

The reduction equation is multiplied by 2 because the number of electrons given in the oxidation

reaction must be equal to the number of electrons gained in the reduction reaction.


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Notice, that pure silver is formed out of the solution of silver salt. Therefore, this process can be

used to produce solid metals out of salts of the same metal.

The above reaction is a general process of replacement. When an element, which is more

reactive electrochemically, reacts with aqueous solution of salt of metal, which is more reactive

electrochemically, displacement reaction will occur.

Another Example: Zinc reacts with copper sulphate, because zinc is more reactive

electrochemically than copper. Therefore, when a strip on zinc in placed in a copper sulphate

solution, the strip is "solvated" in to solution according to the equation:

CuSO4(aq) + Zn(s) Cu(s) + ZnSO4(aq)

Also we can write this reaction into two half reactions

Cu2+(aq) + 2e- Cu (s) (reduction).

Zn(s) Zn2+(aq) + 2e- (oxidation).

Other application of redox reactions is constructing electrochemical cells, which have two types:

Galvanic cells: electrochemical cells in which redox reaction occurs spontaneously and in which

chemical energy converts to electrical energy.

Electrolytic cells: electrochemical cells which uses energy from other source to induce redox

reactions and in which electrical energy converts into chemical energy.

Electrochemical cell composed of two electrodes each of them immersed in a certain electrolytic

solution. Electricity can pass through the cell if

1. Electrodes connected by a conductor (metal wire)

2. Both electrolytic solutions must be connected by a way through which certain

ions are allowed to pass from one solution to another.


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The following figure shows a Galvanic cell in which a total redox reaction occurs.

In this system, an oxidation reaction occurs in beaker1 where atoms of Zn moving out of the

zinc electrode, leaving behind two electrons who travels along the metal wire towards the copper

electrode. Therefore the anode becomes smaller as the redox reaction proceeds.

In beaker2, a reduction reaction occurs where ions of copper are attracted to the cathode, receive

two electrons and precipitate on the electrode as copper atoms. Therefore the cathode becomes

bigger as the redox reaction proceeds.

The salt bridge enables SO42-(aq) to travel between the beakers and close the electric circuit and

retain equilibrium of electric charges in the system.

The redox reaction continues until the system reaches chemical equilibrium.

2 1


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An electrochemical cell is written in the following way:

Zn(s) | Zn2+(aq) || Cu2+ (aq) | Cu(s)

*http://www.sep.alquds.edu/

MATERIALS:

Cu wire, Zn metal, Pb wire, graphite rod, Zn2+ solution, Pb2+ solution, SO42- solution, Fe2+

solution, Fe3+ solution, Cu2+ solution, KCl or KNO3, KmnO4, concentrated H2SO4.

PROCEDURE:

A. Preparation of Cu/Cu2+ Reference electrode

1. Secure a 4-to 5-cm long glass pipet or dropper. Plug the tapered end of the pipet witha small ball of cotton pushed down the pipet with a stif wire. Make sure that the

cotton is tightly packed and that it reaches up to the opening of the pipet.

2. Immerse the tip of the pipet in a beaker containing a saturated solution of KCL or

KNO3 until the cotton is fully soaked.

3. Fill the pipet with 0.1 M CuSO4 solution using another pipet with a rubber bulb.

4. Get a metal rod with a ciameter smaller than the inner diameter of the pipet. Coil a

copper wire around the rod leaving some portion uncoiled for the voltmeter leads.

5. Insert the coiled wire into the pipet. Plug the open portion of the pipet with modeling

clay to fix theposition of the wire. This will also seal the tube and prevent the leakage

of the copper sulfate solution.

Alternative Procedure A.

Anode is always written on

the left side

Cathode is always written on the

right side

Represents the salt

bridge


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1. Secure a 4-to 5- inch long copper wire and dip into a beaker containing 0.1 M

CuSO4 solution.

Prepare a salt bridge as follows: Cut a piece of filter paper into a 1- cm wide, 8-to 10- cm longstrip and dip into a saturated solution of KCl or KNO3. Use a fresh strip for a new solution. DO

NOT USE THE STRIP WITH TWO DIFFERENT SOLUTIONS.


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B. Determination of Electrode Potentials

1. Assemble the following test half-cells:

a. Zn/Zn2+ electrode: Dip a piece of Zn metal strip (approximately 5 cm long, 1 cm wide) in a 50-

mL beaker containing 25 mL of 0.1 M ZnSO4.

b. Pb/Pb2+ electrode: Dip a 10-cm long Pb wire in a 50-mL beaker containing 0.1 M Pb(NO3)2.

c. C/Fe2+, Fe3+ electrode: Dip a graphite rod in a 50-mL beaker which contains 12.5 mL of 0.2 M

Fe(NH4)2(SO4)2 and 12.5 mL of 0.2 M FeNH4(SO4)2.

d. Pb/PbSO4(S), SO42- electrode: Dip a lead wire coated with PbSO4 in a 50 mL beaker which

contains 25 mL of 0.1 M k2SO4. ( To prepare the wire, dip a Pb wire in a solution of 6 M H 2SO4

until the wire becomes coated with PbSO4, a white powdery substance).

2. Set up a cell by coupling a Cu/Cu 2+ electrode with one of the assembled half-cells. If the

reference electrode is that prepared by the first procedure, simply dip the electrode into the

electrolyte of the test electrode. If the second type of Cu/Cu 2+ electrode is used, connect the two

solutions by dipping one end of the filter paper stip in the CuSO4 solution and the other end in

the electrolyt of the test electrode.

3. Measure the voltage of the cell by connecting the Cu/ Cu2+ reference electrode to the positive

(red) terminal of a digital voltmeter and the metal strip or wire or carbon rod of the test electode

to the negative (black) terminal.

4. For the Cu/Cu2+ reference electrode (cathode), the reduction half-reaction and the

corresponding potential is:

Cu2+ + 2 e- Cu(S) = 0.34 volt

The measure potential of the cell is the difference in reduction potentials of the cathode

(reference electrode) and the anode (test electrode).

5. Record your data on the table provided in the data sheet.

C. Electromotive Force of Galvanic Cells


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1. Assemble the following cells by coupling two test electrodes. Connect the two test electrodes

by dipping the ends of a salt bridge (filter paper stip soakeed with saturated KCl or KNO3)

into solutions.

Cell A: Zn/Zn2+ and Pb/Pb2+ electrodes

Cell B: Zn/Zn2+ and C/Fe2+, Fe 3+ electrodes

Cell C: Pb/Pb2+ and C/Fe2+, Fe3+ electrodes

Cell D: Zn/Zn2+ and Pb/PbSO4(S), SO42-electrodes

2. Connect each electrode to the input terminals of the voltmeter and read the voltage. If a

negative voltage is observed, interchange the terminal connections. Note the polarity of the

electrodes. The cathode is the half-cell connected to the positive (red) terminal of the meter,

and the anode is the half-cell connected to the negative (black) terminal.

3. Record the observed cell potentials:

4. Write the cell reactions and determine the theoretical electromotive force of each galvanic cell

using values of electrode potentials from tables.

D. Variation of Electrode Potential with Electrolyte Concentration

1. Prepare several solutions containing varying concentrations of MnO4- and Mn2+ ions. By using

a pipet, measure the following volumes of 0.02 M KmnO4 and 0.05 M Na2C2O4 solution into a

100-mL beaker and mix the resulting solution well.

Mixture mL of KmnO4 mL of Na2C2O4

1 30 5

2 35 10

3 20 15

4 25 20

2. Measure the reduction potential of each mixture using a Cu/Cu2+ reference electrode.

3. Determine the concentration of MnO4- and Mn2+ ions in each mixture.

4. Record the data in the table.

5. Plot cell versus log ([MnO4-]/Mn2+]). Determine the slope of the best straight line.


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6. Based on the slope obtained in step 5, derive an expression relating the cell potential and the

concentrations of MnO4- and Mn2+ ions in the electrode.

Experiment No. 9

GALVANIC CELLS

PRELAB EXERCISES:

NAME: __________________________________________ DATE: _________________

PROFFESOR: ______________________________________

1. Distinguish between galvanic and electrolytic cells.


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2. Distinguish between cell potential and electromotive force.

3. Describe a method for the determination of standard potential of a cell.


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Experiment 9

GALVANIC CELLS


Date: _______________ Score: _______

Group No.: ___________

SIGNATURE

Leader: ____________________________________ __________________________

Members: ____________________________________ __________________________

____________________________________ __________________________

____________________________________ __________________________


OF THE EXPERIMENT)


A. Determination of Electrode Potentials

Test Electrode Reduction Half- Reaction Ecell Eanode

Zn/Zn2+ _____________________________ ________ ________

Pb/Pb2+ _____________________________ ________ ________

C/Fe2+, Fe3+ _____________________________ ________ ________


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Pb/PbSO4(s), SO42- _____________________________ ________ ________


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II. Electromotive Force of Galvanic Cells

Cell Anode Cathode Ecell( observed)

A. Zn/Zn2+ and Pb/Pb2+ ____________ _____________ _______________________

B. Zn/zn2+ and C/Fe2+, Fe3+ ____________ _____________ _______________________

C. Pb/Pb2+ and C/Fe2+, Fe3+ ____________ _____________ _______________________

D. Zn/Zn2+ and Pb/PbSO4(s), SO4____________ _____________ _______________________

Cell A: Reaction 0

Anode : _______________________________________________ ___________

Cathode : _______________________________________________ ___________

Cell : _______________________________________________ ___________

Cell B: Reaction 0

Anode : _______________________________________________ ___________

Cathode : _______________________________________________ ___________

Cell : _______________________________________________ ___________

Cell C: Reaction 0

Anode : _______________________________________________ ___________

Cathode : _______________________________________________ ___________

Cell : _______________________________________________ ___________

Cell D: Reaction 0

Anode : _______________________________________________ ___________


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Cathode : _______________________________________________ ___________

Cell : _______________________________________________ ___________

III. Variation of Electrode Potential with Concentration

Mixture [MnO4]] [Mn2+] log ([MnO4

-]/[Mn2+]) Ecell

1 ________ _________ _________________________

____________

2 ________ _________ _________________________

____________

3 ________ _________ _________________________

____________

4 ________ _________ _________________________

____________

Plot of log ([MnO4]/[Mn2+]) vs. cell


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Slope =__________________ y-intercept =________________

Equation for cell:


1. Does the curve obtained obey the Nernst equation? Why?

2. What is the significance of the slope of the line? How does it compare with the

theoretical value?


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3. What is the significance of the value of the y-intercept obtained in your plot? How does it

compafre with the theoretical value?


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Experiment 10

CONDUCTIMETRY

OBJECTIVES:

1. To be able to determine the relative mobilities of some ions in solution.

2. To be able to determine the variation of conductance with concentration of electrolytes.

3. To be able to determine the concentration of an electrolyte by conductance measurement.

DISCUSSION:

As its name implies, conductimetry is concerned with the electrical conductivity of electrolytes.

Measurements are made indirectly across the resistance of the solution with alternating current,

since direct current would alter the composition of the sample solution by electrolysis.

Experience has shown that for reasons related to the measuring technique (polarization

phenomena) better results are obtained when the measuring frequency is adapted to the range of

measurement.

The conductivity of a solution depends on:

The quantity of ions. The more ions a solution contains, the higher will be its conductivity.

The kind of ions. The smaller and more mobile an ion, the better will be its electrical

conductivity. Thus H3O+, OH-, K+ and Cl- ions all conduct very well.

The solvent. The more polar the solvent, the better the ionisation of the solutes it contains. In

relation to this, water is an ideal solvent, while methyl alcohol is also good.

The temperature. Ionic mobility increases with rising temperature. According to the type of ion,

the conductivity increases by 1-3% / C.

In conductimetric titrations, the cell constant does not usually need to be known, since only the

change in conductance during the titration is tracked.

MATERIALS:

Conductance apparatus (two graphite rods connected in series with a 9-V cell and a digital

multimeter), buret, 0.1M HCl, 0.1 M NaOH, 0.1 M NaCl, 0.1M NH 4C2H3O2.


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PROCEDURE:

A. Determination of Relative Ionic Mobility

1. Place 20 mL of the following solutions in 50-mL beakers:

a. 0.1 M HCl solution

b. 0.1 M NaOH solution

c. 0.1 M NaCl solution

d. 0.1 M NH4Cl solution

e. 0.1 M NaC2H3O2 solution

2. Measure the conductance of each solution by dipping the probes in the solution and

reading from the multimeter the current flowing through the solution. (NOTE: Wash the

probe with distilled water and dry with tissue paper before immersing it into the solution.

3. Tabulate in incrasing order of magnitude the current observed for the solutions containing

the chloride ion. Determine from the results the relative mobility of the cations in the

solutions.

4. Tabulate in increasing order of magnitude the current observed for the solutions

containing the sodium ion. Determine from the results the relative mobility of the anions

in the solutions.

B. Variation of conductance with Concentration

1. Prepare different concentrations of hydro chloric acid by measuring with a pipet the

following volumes of liquid into 50-mL beakers:


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Mixture mL of 1 M HCl mL of distilled water

1 5 20

2 10 15

3 15 10

4 10 15

5 25 0

2. Measure the conductance of each mixture. Calculate the concentration of HCl in the

mixture and record the data in the table.

3. In the absence of extraneous effects, it is expected that the conductance of the

electrolyte is directly proportional to the concentration of the electrolyte (that is, if the

electrolyte concentration is doubled, the conductance of the solution should be

doubled). Based on this premise and on the measured conductance of the most dilute

solution, calculate the expected conductance of each mixture preapared in step 1.

4. Plot the concentration of HCl against the (a)measured conductance and (b) calculated

conductance from step 3 on the same graphing paper.

5. Compare the curves obtained in step 4 and explain the difference between theobserved and expected behavior.

C. Conductimetric Titration

1. Transfer 25.0 mL of the analyte into a 100-mL beaker, dip the probe in the solution

and take the multimeter reading.

2. Slowly add 1 mL increments of the titrant from a buret mixing the reaction mixture

well and recording the multimeter reading after each addition. Stop adding the titrant

when the multimeter reading no longer shows appreciable change in slope. ( This

would be when appoximately 35 mL of titrant has been added.)

3. Tabulate the meter readings and the total volume of titrant added.


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4. Plot the meter reading against the total volume of titrant added.

5. Determine from the titration curve the volume of the titrant needed to reach the

equivalence point.

6. Calculate the concentration of the analyte.


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Experiment No. 10

CONDUCTIMETRY

PRELAB EXERCISES:

NAME: __________________________________________ DATE: _________________

PROFFESOR: ______________________________________

1. Discuss the following:

a. Conductance and Conductivity

b. Strong and Weak Electrolytes

c. Mobilities of ions and Conductivity


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Experiment 10

CONDUCTIMETRY


Date: _______________ Score: _______

Group No.: ___________

SIGNATURE

Leader: ____________________________________ __________________________

Members: ____________________________________ __________________________

____________________________________ __________________________

____________________________________ __________________________


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OF THE EXPERIMENT)


A1. Determination of Relative Ionic Mobility

Solution I (mA)

Hydrochloric acid __________

Sodium hydroxide __________

Sodium chloride __________

Ammonium chloride __________

Sodium acetate __________


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A2. Relative Mobility of Cations

Cation I(mA) Conclusion

H+ _________________

________________________

Na+ _________________

________________________

NH4+ _________________

________________________

A3. Relative Mobility of Anions

Anion I(mA) Conclusion

OH- ___________________________ ____________________________________

Cl- __________________

________________________

C2H3O2- __________________

________________________


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Relative Mobility:

B. Variation of Conductance with Concentration

Mixture [HCl]mixture I(mA)

mL of 1 M HCl mL of distilled H2O Measured Expected

5 20 ______________ _________

____________

10 15 ______________ _________

____________

15 10 ______________ _________

____________

20 5 ______________ _________

____________

25 0 ______________ _________

____________


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Plot of [HCl] vs. Conductance (mA)

Compare the curves (a) and (b) and explain the difference between observed and expected

bahavior.

C. Conductimetric Titration

Vtitrant I(mA) Vtitrant I(mA) Vtitrant I(mA)


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______ ______ ______ ______ ______ ______

______ ______ ______ ______ ______ ______

______ ______ ______ ______ ______ ______

______ ______ ______ ______ ______ ______

______ ______ ______ ______ ______ ______

Conductimetric titration Curve (Plot of Vtitrant vs. Conductance)


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Vtitrant at equivalent point = _________________________

Concentration of Unknown =__________________________

phy chem 2 lab

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