Course Manual 13-14

8/9/2019 Course Manual 13-14

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NOTICE TO ALL CHEMISTRY STUDENTS

All students working in the Chemistry undergraduate laboratories MUST WEAR INDIRECTLY-VENTED CHEMICAL SPLASH SAFETY GOGGLES AT ALL TIMES.

All students must familiarize themselves with the safety rules pertaining to a particular experiment(e.g. use fume hood, wear gloves, do not pipet by mouth, etc.).

STUDENTS WITH ANY KIND OF MEDICALLY RELATED CONDITIONS (e.g. seizures,etc.) or who are pregnant must contact the course instructor for information regarding the advisabilityof taking this lab course.

ANY STUDENT WHO NEGLECTS TO FOLLOW THE SPECIFIED, SAFE PROCEDUREFOR CONDUCTING AN EXPERIMENT AS DESCRIBED IN THIS MANUAL, OR WHOPERSISTS IN REFUSING TO FOLLOW THE NORMAL SAFETY RULES, WILL BEASKED TO LEAVE THE LABORATORY FOR THE REMAINDER OF THELABORATORY PERIOD.


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TABLE OF CONTENTS

Introduction 1

Lecture and Test Schedule 2

Laboratory and Tutorial Schedule 3

Textbook Assignments (for practice) 4

Additional Problems - Unit 1 - Solution Concentrations (for practice) 5

Tutorial Assignments 1 - 5 (for credit) 6

Lecture Note Supplement - Unit 2 - Solution Equilibria 12

Marking Scheme 30

Policies Regarding Work Not Done or Submitted Late 30

General Laboratory Instructions 31

Materials 33

Safety 33

Laboratory Practice 34

The Laboratory Notebook 36

The Report 37

Laboratory Report Marking Scheme 39

Integrity and Ethics in the Laboratory 39

Experiments

1. Separation of Metal Ions by Paper Chromatography 41

2. Behaviour of Gases 53


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3. Analysis of Antacids by Acid-Base Titration 65

4. Spectrophotometric Determination of the Formation Constant of a 75Complex Ion

5. Measurement of the Enthalpy of Reaction by Calorimetry 85

Appendices

A. Error Analysis 103

B. Beer-Lambert Law 110

C. Spectronic 20114


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CHM 110H5F

Chemical Principles 12013 - 2014

http://www.utm.utoronto.ca/~w3chm140

Welcome to CHM110H. I hope that through this course you will come to appreciate chemistry asan integral part of our culture. You should also become comfortable in dealing with a changing bodyof knowledge through developing an appreciation of the processes of chemical research. I look forward to meeting each of you, to advising on your programs in the sciences, and to mentoring youas you progress towards your goals.

Instructor: Judith PoRoom: 4048, Davis Building

Phone: (905) 828-3803E-mail: preferrably via Virtual Office Hours on the course website listedabove or, if VOH is not available, at [email protected]

Office Hours: M, W 12:15-1:30p.m.

LectureTimes: M, W, F 9-10 Room IB110

orM, W, F 11-12 Room KN137

Texts: M.S. Silberberg, S. Lavieri and R. Venkataswaran, Chemistry: TheMolecular Nature of Matter and Change, Canadian Edition, McGraw-HillRyerson (2013) - with Students Solutions Manual and Connect

CHM110H Course Manual - You must print this, put it in a binder andbring it to each laboratory and tutorial class.

OtherRequiredMaterials: 1. indirectly-vented, chemical splash safety goggles

2. lab coat (100% cotton recommended)3. disposable gloves and/or kitchen type gloves4. non-programmable calculator (Note that the only calculators that

will be allowed in tests and exams are the following: TI-30XIISor CASIO fx-260 solar) .
http://www.utm.utoronto.ca/~w3chm140mailto:[email protected]:[email protected]://www.utm.utoronto.ca/~w3chm140


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Lecture and Test Schedule 2013-2014

StartingDate Unit Numberof Lectures

Chapters Topic

Sept. 9 1 Introduction to the Course and itsWeb Site

Sept. 11 1 8 1 - 4 and 12.3-12.5

4

Matter, Reactions and SolutionStoichiometry

Behaviour of Gases

Sept. 30 2 11 15 - 17 Equilibria

Oct. 30 3 8

5

5 and 18

19.1-19.6

Thermodynamics

Electrochemistry

.

Mid-term Tests: Monday, October 7, 8:00-9:00 a.m. (no CHM110H lectures on this day)

Monday, November 4, 8:00-9:00a.m.(no CHM110H lectures on this day)

Students with another class at this time must inform the instructor by e-mail of the course and the room in which it meets at least one week inadvance of the test. Those students only will be allowed to write the testsfrom 9-10 a.m.

Final Exam: Sometime in the period of December 9 - 20, time and place to be determined

Personal plans for this time period that interfere with your availabilityto write a final exam are not considered legitimate excuses for missingan exam. Therefore do not make any personal plans for this time perioduntil the exam schedule is published by the Registrars Office.


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Laboratory and Tutorial Schedule 2013-2014

Week beginning

L/T Work to be Done Work Due

Sept. 9 L Laboratory Safety, Exercises in the Use of Laboratory Equipment

16 T Assignment 1Quiz 1

23 L Exp. 1 - Separation of Metal Ions by Paper Chromatography

A. 1

30 T Assignment 2Quiz 2 Exp. 1

Oct. 7 L Exp. 2 - Behaviour of Gases A. 2

14 T Assignment 3 Exp. 2

21 L Exp. 3 - Analysis of Antacids by Acid-BaseTitration

A. 3


Exp. 3

Nov. 4 L Exp. 4 - Spectrophotometric Determination of the Formation Constant of a Complex Ion

A. 4


Exp. 4

18 T Quiz 5 A. 5

25 L Exp. 5 - Measurement of Enthalpy of Reaction by Calorimetry

Exp. 5 *

All lab reports and tutorial assignments must be submitted in hard copy.

* Note that lab reports are due in your tutorial class in the weeks noted in the schedule withthe exception of Exp. 5. All reports for Exp. 5 are due by 5:00p.m. on Wednesday, December4..


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Textbook Assignments

The following Chapters will be studied in the fall term:Unit 1 - Stoichiometry - 1-4 and 12.3-12.5Unit 2 - Equilibria - 15-17 (plus notes on pages 12-29 of this manual)Unit 3 - Thermodynamics and Electrochemistry - 5, 18 and 19.1-19.6

The assigned problems are listed below. It is recommended that you focus on the problemsnumbered in green whose answers are in Appendix D and whose complete solutions are in theStudent Solutions Manual. These problems will be discussed in your tutorials. These problems arenot to be handed in, however their content will be reflected in the quizzes.

UNIT 1Chapters 1 - 4

all green problems

Chapter 12 problems 12.44 - 12.70 plus additional problems on page 5 of this manual

UNIT 2Chapter 15 - 17

all green problems

UNIT 3Chapter 5 and 18

all green problems

Chapter 19green problems 1 - 81 and 121, 125, 137, 154


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Additional Problems - Unit 1Solution Concentrations

Consider each of the following aqueous solutions at 25 oC.

1. Given: H 2SO 4 98.0% w/w = 1.842 g mL-1

Find: molarity, molality and mole fraction of acid

2. Given: Cr 2(SO 4)3 1.26 M 1.37 m

Find: mole fraction, density and % w/w

3. Given: H 3PO 4 = 1.412 g mL-1

Xacid = 0.510

Find: molarity, molality and % w/w

4. Given: C 2H4(OH) 2 4.028 m = 1.024 g mL-1

ethylene glycol

Find: molarity, % w/w and mole fraction of ethylene glycol

Answers

1. 18.4 M, 500 m, X = 0.900

2. X = 0.0241, = 1.41 g mL -1, 35.0 % w/w

3. 57.8 m, 12.2 M, 85.0 % w/w

4. 3.299 M, 19.97 % w/w, X = 0.06756


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Tutorial AssignmentsGeneral Instructions

Unlike the other assignments, tutorial assignments are to be handed in and will be marked. Youshould attempt to do the assignment before the tutorial in which it will be considered. In the tutorial,the TA will not tell you how to do the assignment. But he/she will guide you in an approach to theassignment and will try to answer your questions while not telling you the solution. You will thenhave an opportunity to revise and improve upon your initial work before handing in the assignmentthe following week. To gain the most benefit from the tutorial and the opportunity to revise your work, you must have come to the tutorial prepared, i.e. having already worked on the assignment.

Formatting:

1. Do not include a cover page.

2. The following information should appear cross the top of the first page:

name and student number, lab section number, TAs name, assignment number

3. Use 1 inch margins, 1.5 line spacing and 12 point Times New Roman font (10 point for references).

4. Assignments must not exceed two pages in length, including any graphs, calculations andreferences where appropriate.

5. References should be given superscript numbers in the text and listed at the end of theassignment.

Assignments are due in your laboratory or tutorial class in the weeks noted in the schedule on page3 of this manual. The penalty for late assignments is 5% off of the assignment mark per calendar dayto a maximum of 7 days after which a mark of zero will be given.

Unlike experiments, quizzes and tests which can only be done on a particular day, assignments can be done over a period of many days or weeks. Therefore medical or other excuses will not beaccepted as a reason for missing an assignment (excepting, of course, in the unfortunate circumstanceof a prolonged, serious illness).


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Assignment 1

a. A 1.00 g sample of enriched water, a mixture of H 2O and D 2O, reacted completely with Cl 2

to give a mixture of HCl and DCl. The HCl and DCl were then dissolved in pure H 2O to

make a 1.00 L solution. A 25.00 mL sample of the 1.00 L solution was reacted with excess

AgNO 3 and 0.3800 g of an AgCl precipitate formed. What was the mass % of D 2O in the

original sample of enriched water?

Atomic masses (g/mol): H = 1.0, D = 2.0, O = 16.0, Cl = 35.5, Ag = 107.9

b. The mass % natural abundance of the isotopes H and D are 99.985 and 0.015 respectively.

The mass % natural abundance of the isotopes16

O,17

O and18

O are 99.759, 0.037 and 0.204respectively. In a mass spectrometry experiment on naturally occurring water, at what m/

values would you expect to observe the three most abundant peaks for the molecular ions and

what would you expect to be their relative abundances?

c. For questions a. and b. above, explain the process by which you arrived at a plausible solution.

In developing your answer, justify the steps you have undertaken (i.e., use analogies/examples

to support the choice of steps that you have undertaken in solving the two questions above).


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Assignment 2

P (atm) 1.00000 0.66667 0.50000 0.33333 0.25000 0.16667

PV (L atm) 22.2643 22.3148 22.3397 22.3654 22.3775 22.3897

Data for PV as a function of P for 1 mol of CO 2 at 0 C is given in the Table above.

a. Plot PV as a function of P on a scale sufficiently expanded so that the experimental variationsin PV can be observed on the graph.

b. From the plot, determine the value of RT at 0 C.

c. The plot follows the equation, PV = RT + BP where B is an empirical constant. Determinethe value of B.

d. On the same sheet plot PV versus P for an ideal gas.

e. Calculate the value of PV at 0.90000 atm for 1 mol of CO 2 at 0 C using the equationdetermined in c. above. What would be the percentage error made if you used the ideal gaslaw to calculate PV at 0.90000 atm for 1 mol of CO 2 at 0 C ?

f. Repeat the calculation in e. but at P = 90.000 atm. Comment on any difference in the percentage error for the two calculations.


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Assignment 3

Captain Kirk, of the Starship Enterprise, has been told by his superiors that only a chemist can be

trusted with the combination to the safe containing the dilithium crystals that power the ship. The

combination is the pH of solution A described below, followed by the pH of solution C. (Example:

If the pH of solution A is 3.47 and that of solution C is 8.15, then the combination to the safe is 3-47-

8-15.) The chemist must determine the combination using only the information below (all solutions

are at 25C).

Solution A is 50.0 mL of a 0.100 M solution of the weak monoprotic acid, HX.

Solution B is a 0.0500 M solution of the salt NaX. It has a pH of 10.02.

Solution C is made by adding 15.0 mL of 0.250 M KOH to solution A.

What is the combination to the safe?

If, in preparing solution C, 15.0 mL of water was added instead of the 15.0 mL of KOH, what effect

would this have on the combination to the safe?

Show all calculations.


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Assignment 4

The municipal water processing plants of communities with hard water often treat the water supply

with slaked lime, Ca(OH) 2 , in order to remove Ca 2+ from the water. The slaked lime reacts with bicarbonate from the dissolved metal bicarbonates in the water according to the following reaction:

Ca(OH) 2 + 2 HCO 3 CaCO 3 (s) + CO 32 + H 2O .

The carbonate produced then reacts with Ca 2+ originally in the water to precipitate as CaCO 3 (s).Thus calcium ions from slaked lime are added to the hard water in order to remove calcium ions fromthe hard water!

If hard water has a calcium ion concentration of 1.8 x 10 3 M and you want to reduce thatconcentration to 0.6 x 10 3 M, what must be the concentration of the slaked lime? If you didnt dothis calculation and you simply used a saturated solution of slaked lime, what do you think theconsequences might be? Support or refute your prediction by quantitatively assessing the likelihoodof your predicted consequences.


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Assignment 5

a. The sun supplies energy at a rate of about 1.0 kilowatt per square meter of surface area. The

plants in an agricultural field produce the equivalent of 20.0 kg of sucrose per hour per hectare

(1 ha = 10,000 m 2 ). Assuming that sucrose is produced by the reaction

12 CO 2 (g) + 11 H 2O (l) C12H22O11 (s) + 12 O 2 (g) H = 5640 kJ ,

calculate the percentage of sunlight used to produce the sucrose, i.e. determine the efficiency

of photosynthesis.

b. The best solar panels currently available are about 15% efficient in converting sunlight to

electricity. A typical home will use about 40 kWh of electricity per day (kWh = kilowatt

hour). Assuming 8.0 hours of useful sunlight per day, calculate the minimum solar panel

surface area necessary to provide all of a typical homes electricity.

c. Describe the methodology you followed in determining the minimum solar panel surface area

in question b. above. In your response provide a justification for each step you undertook in

solving the problem.


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Lecture Note Supplement - Unit 2

Solution Equilibria

The following procedure allows one to deal quantitatively with any chemical system at equilibrium.

1. Write balanced chemical equations to represent all of the reactions that occur in the system.

2. List all species that exist in the system at equilibrium (omitting the solvent if it is present inlarge excess).

3. Write a number of independent, simultaneous equations, equal in number to the number of species in the system, which relate the equilibrium constants for the reactions occurring andthe concentrations of the species in the solution. These equations will be of three types:

a. equilibrium constant expressions, b. mass balance equations andc. a charge balance equation.

4. Assess the values of equilibrium constants and concentrations that are known to determineif it is likely that the concentration of any one species is negligible as compared to that of another species. If this is so, it may be possible to simplify the mass balance and/or the charge

balance equations.

5. After simplifying the equations where possible, solve the series of simultaneous equations for the concentration or equilibrium constant that is being sought.

6. Based upon the solution, check to ensure that any simplifying assumptions that were madewere valid.

7. If the simplifications were justified, the problem is then successfully completed; if they werenot, then the full set of simultaneous equations must be solved.

A system at equilibrium can be quantitatively described by the following properties: C, the originalconcentration of each reactant; V, the volume of each reactant; K, the equilibrium constant for eachreaction; and [species], the equilibrium concentration of each species present. If some of these

properties are known or measured experimentally, the others can be calculated by using the procedure

described above. For example, if C, V and K are known, the equilibrium concentration of all species present can be calculated.

On the following pages are steps 1-3 of the general treatments of some common types of chemicalsystems involving acids, bases and salts in aqueous solution.


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Marking Scheme

Experiments: (5% each) 25

Quizzes (1.5% each) 7.5

Tutorial Assignments (1.5% each) 7.5

1 Hour Test - October 7, 2013 12.5

1 Hour Test - November 4, 2013 12.5

Final Exam - TBA, December 9 - 20, 2011 35

Total 100

All marks represent % of Total mark.

Policies Regarding Marks for Work Not Done or Submitted LateStudents are required to declare their absence from a class for any reason through their ROSIaccounts in order to receive academic accommodation for any course work such as missed tests, lateassignments, and final examinations. Absences include those due to illness, death in the family,religious accommodation or other circumstances beyond their control In addition, students mustfollow the instructions below.

1. Experiments, Quizzes and 1 Hour Tests

All absences must be declared on ROSI . In addition, within one week of the date of themissed work, students should submit to the course instructor a signed letter explaining the

reason for their absence. The letter should include the students name, phone number, e-mailaddress, student number and lab section number as well as the date of and the description of the missed work. For absence due to illness, an official U of T Medical Certificate isrequired. That Certificate or other documentation appropriate to the reason for the absenceshould be stapled to the letter. If the explanation is deemed reasonable (after thedocumentation is verified), the final exam mark will be used as the mark for the missedwork. If the explanation is unreasonable or if no letter is submitted within one week of themissed work, a mark of zero will be given for the missed work.

THERE WILL BE NO MAKE-UP EXPERIMENTS, QUIZZES OR 1 HOUR TESTS.

2. AssignmentsUnlike experiments and tests which can only be done on a particular day, assignments can

be done over a period of many days or weeks. Therefore medical or other excuses will not be accepted as a reason for missing an assignment (excepting, of course, in the unfortunatecircumstance of a prolonged, serious illness).

3. The penalty for late submission of an assignment or laboratory report is 5 marks off per calendar day to a maximum of 7 days, after which a mark of zero will be given.

4. Final Examination: Refer to the UTM Academic Calendar for these regulations.


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General Laboratory Instructions

We should first of all consider why laboratory experience is an essential part of any university

science program. Why is it very important to learn to make careful observations and experimental

measurements in the laboratory?

All of our accumulated understanding of the physical world ultimately depends on the large

number of careful experimental observations that have been made by many generations of scientists

since the beginning of recorded history, and to which research scientists are adding every day. In

chemistry it is convenient to divide experimental observations into two main categories: those that

we refer to as qualitative observations and those that we refer to as quantitative observations.

Qualitative experiments involve observations with our normal physical senses, such as sight and

smell. For example, the observation of a color change in a chemical reaction is a qualitative

observation. Quantitative observations involve the measurement of some physical quantity such as

mass, volume, pressure, concentration or temperature. Whether we are making a qualitative

observation or a quantitative measurement in the laboratory, it must be done as carefully as possible

and reported with complete honesty.

It is of the greatest importance that all of the observations that you make during this

laboratory course be recorded immediately in this manual which also serves as your

laboratory notebook. If you are doing a qualitative experiment, immediately describe as carefully

and as accurately as you can what you actually observe, no matter what you might have anticipated

on the basis of previous knowledge or theoretical considerations. If you are making a quantitative

measurement, immediately record any quantity that you measure with its correct units. If you think

that the result is strange or unusual, say so, but do not tamper with the results. There is no place for

fiction in the chemistry laboratory. There is no such thing as an incorrect experimental result. Many

new discoveries have resulted from experiments that went wrong or did not fit the theory.


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Honest mistakes can of course occur and inaccuracies may result from lack of care or imprecision

in making measurements. That is why we repeat experiments when we can, so that we have a basis

for deciding what results might legitimately be discarded. When you get an unusual or imprecise

result you should always try to understand why. What was there about the experimental procedure

that might have led to error? Consult your TA if you are puzzled about a result or think that you may

have carried out some procedure inaccurately.

The final objective of many experiments is simply to support our understanding of well

understood concepts and theories. In others we may be interested in improving an experimental

procedure to get a more accurate result or to improve the yield of a compound that we are making.

If you can think of ways in which an experiment could be improved, say so. In general, however,

especially in university, an important objective is to use new experimental observations as the basis

for improving existing theories or formulating new concepts that will ultimately lead to better

theories that embrace more experimental observations than the old theories. The results of careful

experiments have a timeless quality; they are, indeed, the permanent body of knowledge that

constitutes science. Scientists may repeat your experiment in the future; if it is accurate they will

still get the same result. The only difference might be that they can improve the accuracy of a

physical measurement because they have improved measuring instruments. In contrast, theories are

simply the best models that we can formulate to tie together as many as possible of the facts that are

known today. As new results accumulate old theories are discarded and replaced by new ones.

The object of a laboratory course, therefore, is not simply to mindlessly get correct answers

but to learn to understand how to handle common laboratory apparatus and what needs to be done

to get reliable results. In terms of rewards, both short-term and long-term, a careful approach and

an aware, enquiring mind count for far more than a slavish attempt to reproduce what the instructions

or textbooks say is right.


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1. Materials

YOU WILL NOT BE ALLOWED TO WORK IN THE LABORATORY WITHOUT

ALL OF THESE MATERIALS !!

a. This Course Manual/Notebook (the entire manual, not just selected pages) must be

brought to all laboratory classes.

b. A laboratory coat (100 % cotton recommended although other materials are

acceptable) to protect you and your clothes is mandatory .

c. Indirectly vented chemical splash safety goggles ( on sale in the Bookstore) must be

worn at all times in the laboratory. Gloves are required for some procedures.

2. Safety Precautions

Experimental chemistry is inherently dangerous; many experiments can be hazardous unless

the experimentalist is aware of the nature of the materials used and is careful in handling

them. The best precaution against accident is to understand what you are doing and to keep

a neat and well-organized laboratory bench.

THE FOLLOWING PRECAUTIONS MUST BE OBSERVED.

a. Eye protection must be worn at all times. Indirectly vented chemical splash safety

goggles are on sale in the Bookstore. (N.B. Concentrated alkalies, such as 30%

sodium hydroxide, can dissolve the cornea instantaneously.)

b. If a chemical accidentally gets in your eyes, in your mouth, or on your skin, rinse the

affected area immediately with plenty of cold water. Do not delay in doing this,

whether it involves you or a neighbor. Immediate action can prevent a serious injury.

Then report the accident to your TA or the technician, who will decide if further

treatment is needed.

c. Never taste a chemical. Consider all chemicals as potentially toxic. Always wash


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your hands before leaving the laboratory.

d. Clean up chemical spills immediately. For acid or base spills, rinse off with water

and inform the TA or technician.

e. Read the labels on all reagent bottles carefully. Be sure that you know what

chemical you require and in what quantity. If it is a solution, carefully check the

label to make sure it is in the correct concentration. Serious hazards can result from

mixing (mistakenly) certain solutions.

f. Never pipet by mouth. Your TA will show you how to measure out fixed volumes

of solutions using a pump or a rubber bulb on the pipet.

g. Perform the experiment in a fume hood if corrosive or toxic vapors are in any way

involved in the experiment.

h. Dispose of waste materials properly.

(1) Broken glass should be put only in the containers marked Glass.

(2) Waste chemical solids should be disposed of in the special containers

provided. Do not mix waste chemicals and waste paper.

(3) Check with your TA before disposing of any waste liquids down the sink.

For certain liquids, special waste containers will be provided.

i. Eating and drinking are strictly forbidden in the laboratory .

j. Dress appropriately. Do not wear sandals or shorts. Tie back long hair. A lab coat

is mandatory .

3. Laboratory Practice

a. Reagents. These are usually obtained from stock bottles. If you contaminate the

stock bottle or remove unnecessarily large amounts of a reagent, then you have


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committed an anti-social act.

(1) Never remove reagent bottles from the supply area.

(2) When using stock reagents, keep the bottle stoppers clean and always replace

them after use.

(3) Know how much reagent you require and take only the amount needed.

(1) Use a clean, dry spatula in handling solids.

(2) Never insert a pipet into a reagent bottle. Instead, transfer the necessary

solution to a clean, dry beaker and pipet from the beaker.

(3) Never return unused chemicals to the stock reagent bottles

b. Balances. Two kinds of balances are available. For the most accurate weighings,

use an analytical balance. For weighings that require less accuracy, use a triple-beam

balance. Make sure that you know which balance is required for a particular

weighing. Under no conditions should reagents be allowed to come into contact

with the balance pans. When using the analytical balance weigh by difference from

a clean, dry weighing bottle. If you have an accidental spill, report it immediately to

your TA or technician. When using the triple-beam balance, weigh by difference into

a clean beaker or onto a clean piece of weighing paper.

c. Distilled Water. The supply of distilled water is limited. Use it only for the final

rinsing of glassware. It should also be used for making up all aqueous solutions. The

use of ordinary tap water can lead to spurious results.

d. Experimental Hints

(1) Cleanliness is essential. Clean all glassware and rinse with distilled water.

(2) For reactions which need to be carried out at elevated temperatures, a hot


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plate or a hot water bath on a hot plate should be used.

(3) In separating a precipitate from a solution by centrifugation, be sure that the

centrifuge is properly balanced.

(4) Whenever reagents are combined, be sure that the resulting solution is

thoroughly mixed. Diffusion in solution can be a very slow process.

4. The Laboratory Notebook

Laboratory work may not be done without this Manual (used also as a Notebook).

This notebook is your most important piece of equipment in the laboratory, so important that

you should never be in the laboratory without it. It is used to keep a record of all the details

and observations of the experiments performed. The following guidelines should be

observed.

a. Record all entries in permanent, water-proof ink. If for some reason you want an

entry to be disregarded, it should be crossed out (not erased) and the reason for

disregarding it should be noted

b. Before coming to the laboratory, read the experiment carefully. Answer all the pre-

lab questions for the experiment. Pre-Lab Questions must be handed in to your

TA as soon as you enter the lab. This will be followed by a 5 minute lab quiz.

The quiz question will be one of the pre-lab questions. This must be written and

handed in before you begin any experimental work .

c. Any procedures used, if they differ from the instructions given and any observations

made should be recorded in the notebook on the back of the data sheet as the

experiment is being performed. This can be done in note form. If you are using a

balance, take your notebook with you to the balance to record your data. Trying to


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remember afterwards or jotting down observations on scraps of paper is unreliable

and represents poor laboratory technique.

d. Before leaving the laboratory, have your pages of data initialed by your demonstrator.

5. The Report

Reports should contain the following items:

a. a Cover Page containing your name and student number, the full name of your lab

partner if you worked in pairs, your lab section number, your TAs name; the number

and title of the experiment, the date on which it was performed and the date on which

the report was submitted;

b. a few lines, with chemical equations where applicable, describing the Purpose of the

experiment;

c. the Experimental Method - the actual record of how the experiment was done,

taken from your notebook record. Deviations from the procedure given in the

manual should be described in detail. The instructions in this manual should not

be mindlessly rewritten in your report, but should be referenced. Therefore this

section may be very brief unless there were significant deviations from the procedure

in the Course Manual.;

d. the experimental Results including observations, tables of data, calculations and

graphs where appropriate and answers to any questions in the text of the experiment;

e. a brief discussion of the results which may include comparison with theory,

quantitative assessment of precision and accuracy, explanation of errors, or

suggestions for improvement of the experiment;

f. a Summary of the results and conclusions, given in a few lines (or in a table if


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appropriate);

g. References - the source of any information written in your report that was obtained

from books, journals, websites or other sources:

(1) booksAuthor Title ; Publisher : PlaceZumdahl, S. Chemical Principles , 6 th edition; Brooks/Cole.: Belmont,CA;Year Page 2009; p 47

(2) journalsAuthor Journal Year , Vol., PageAggarwal, V.A. J. Am. Chem. Soc. 1996 , 118, 7004

(3) websiteslast

Author Title URL updateHsu, D. Chemicool Periodic Table ; http://www.chemicool.com; Aug. 5,2013

h. The Data Sheets from this Manual should be attached at the end of the report.

i. For Experiments 1, 3 and 4, the report should also address the Writing Initiative

assignment for that experiment (found immediately before the pre-lab questions for

that experiment). This part of the report should be a maximum of one page, with 1.5

line spacing, 12 point Times New Roman font, and 3/4 inch margins. Your name,

student number, PRA section number and TAs name should be printed in a single

line at the top of the page. This should be handed in at the same time as the rest of the

report but should not be stapled to it.

All reports must be prepared with the use of word processing and spreadsheet programs

If you are not familiar with the use of spreadsheets, visit the following website for

instruction: http://library.utm.utoronto.ca/excel .

Reports are due as listed in the Schedule on page 3.
http://library.utm.utoronto.ca/excelhttp://library.utm.utoronto.ca/excel


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7. Marking Scheme for Laboratory Reports

The nature of each experiment is different and therefore each will have a different marking

scheme. For example, in some qualitative observations are most important, in others,

quantitative calculations and graphs are prominent. However 10% of each report mark will

be for report presentation. This refers to formatting, appearance, grammar, spelling, clarity,

and following the report guidelines given above. The writing assignment in the reports for

Experiments 1, 3 and 4 will each count for 15% of the total report mark. Another 10% of

each report mark will be for experimental technique. This refers to your actual performance

in the laboratory. The table below lists some common ways that technique marks are lost.

As you can see they are easy to avoid.

Infraction Mark Deduction Infraction Mark Deduction

not bringing goggles* -2 not recording data in permanent ink

-3

not bringing lab coat -2not readingequipment to itstolerance level

-3

leaving a messaround the balances

-2 (for everyone inthe section)

leaving a dirtyworkstation at theend of the lab period

-2

accidentally breakingglassware**

0

*Students are not allowed in the lab without safety goggles. There is a limited supply of goggles thatmay be borrowed from the lab technician, at a cost of 2 marks. If this supply runs out and you donthave goggles, you will not be allowed in the lab and will get zero for that experiment mark..

**NB While there are no deductions for accidently breaking glassware, this is contingent on youreporting any breakages to your TA. Not reporting broken glassware is a major safety violation andwill carry a significant mark deduction.

8. Integrity and Ethics in the Laboratory

A scientists most valuable possession is integrity. Be a scientist! Be conscientious

in your efforts to observe, collect, record, and interpret the experimental data as best as you


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can. In CHM110H and at the University of Toronto Mississauga, only honest scientific work

is acceptable.

Occasionally you make a measurement that you think is incorrect. At such a time you

may be tempted to change the measurement or to copy the measurement of another student.

I urge you to strongly resist this temptation. A person who alters their data is of no use in

the scientific community. As well, the academic penalties for such behaviour are severe, the

minimum penalty being a mark of zero for that experiment and the notation of an academic

offence on your official academic record.

From the Code of Behaviour on Academic Matters:

It shall be an offence for a student knowingly:

(d) to represent as ones own any idea or expression of an idea or work of another in

any academic examination or term test or in connection with any other form of

academic work, i.e. to commit plagiarism.

The quotation above is taken from an excellent article written by Senior Lecturer Emeritus,

Margaret Procter, entitled How Not To Plagiarize. It is well worth reading and can be

found at the following address:

http://www.writing.utoronto.ca/advice/using-sources/how-not-to-plagiarize


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EXPERIMENT 1

Separation of Metal Ions by Paper Chromatography

Introduction

The first use of chromatography was described by the Russian botanist, Mikhail Tsvet, in

1906. He used a solvent to carry coloured material extracted from vegetables along a length of paper

and demonstrated the separation of - and -carotene. Tsvet named the process chromatography, from

the Greek for colour writing (chromos and graph). Today the term applies to a number of methods

for separating the components of a mixture on a support material.

In paper chromatography, a sample spot of the mixture to be analyzed is applied on to an

adsorbent paper (the stationary phase). The end of the paper is then dipped into a solvent (the mobile

phase) which then rises up the paper by capillary action. As the solvent passes over the sample spot,

the components of the mixture are attracted to the solvent and are carried with it up the paper. But

the rate at which each component of the mixture is carried up depends upon how strongly the

component is attracted to the paper as compared to how strongly it interacts with the solvent. As a

result, each component of the mixture rises up to a different position on the paper thus separating the

components of the mixture. The ratio of the distance that a compound rises up the paper to the

distance that the solvent moves up the paper is called the retention factor, R f . The R f values for

compounds are dependent on the solvent and the temperature.. If the components of the mixture are

coloured, their positions on the paper can be seen by eye. If they are colourless, they can be detected

by viewing the paper under UV light or by exposing the paper to another compound that reacts with

the components to form coloured compounds.

In this experiment the R f values for five different transition metal cations will be determined.

In addition, an unknown mixture of some of these cations will be analyzed.


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Pre-lab Questions

Answer the pre-lab questions found on pages 49-50 . Those pages must be removed from this

manual and submitted to your TA when you enter the laboratory.

Procedure

In this procedure you will apply eight sample spots to a piece of chromatography paper: one spot of

each of the five metal ion containing solutions, one spot of a solution that contains all five of the

metal ions, and two spots of your unknown solution (which will contain from 2 to 4 metal ions).

After the spots dry, the paper will be dipped into the solvent in the developing chamber and the

chromatogram will be allowed to develop. When the solvent front has reached up to within about 1.5-

1.0 cm of the top of the paper, the paper will be removed from the developing chamber and the

solvent front will be marked with a pencil. When the paper is dry, any visible bands will be circled.

The chromatogram will then be enhanced by being placed in the ammonia chamber for 5-10 minutes

and any additional bands that appear will be circled. Finally, the bands will be further enhanced by

reaction with the reagents listed in the Table. Note that gloves must be worn for all work with the

developing solvent and the ammonia chamber and this work MUST be done in the fume hood.

Gloves should also be worn when handling the chromatography paper.

1. Preparation of the Developing Chamber

Using a graduated cylinder, measure out approximately 10mL of the developing solvent (a 9:1

mixture of acetone:6M HCl). Pour the solvent down a glass stirring rod into the middle of

a dry 600mL beaker . It is important that you do not wet the sides of the beaker. It is also

critically important that the height of the liquid in the beaker is less than 1cm. Cover

the beaker tightly with plastic wrap in order to allow the atmosphere in the beaker to become

saturated with the vapour of the solvent. (Saturation takes about 10 minutes.)

2. Preparation of the Ammonia Chamber


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The ammonia chamber consists of an 800/1000mL beaker in the center of which is a 30/50mL

beaker which contains about 5mL of concentrated aqueous ammonia solution. The larger

beaker is covered tightly with plastic wrap. This chamber will have been preassembled for

you and will be in your fume hood. Check to see that there is about 5mL of ammonia solution

in the small beaker. If there is not, ask your TA or the technician for asistance in topping up

the ammonia level. This chamber must be kept in the fume hood and must be tightly

covered with plastic wrap at all times excepting when placing or removing the

chromatography paper.

Figure 1. Preparation of the Chromatography Paper

3. Preparation of the Chromatography Paper

Wearing gloves , obtain one 10cm x 20cm chromatography paper and place it on a clean piece

of ordinary white paper, not directly on the lab bench. Always wear gloves when handling

the chromatography paper. With a pencil, measure and mark the chromatography paper as

shown in Figure 1 above. Where the tick marks cross the 1.5cm line is where you will apply

the sample spots


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4. Applying the Sample Spots

The liquid sample spots will be applied to the chromatography paper using a capillary

tube. The best chromatograms result from small, compact sample spots. Aim for a spot size

between 2 and 4 mm in diameter. You may wish to practice by spotting distilled water along

the top 20cm edge of the paper, confining the water spots to within 1cm of the top of the paper.

Once you are confident of your spotting skill, proceed to apply the sample spots.

You will find six labelled solutions in test tubes on your bench, five of which contain

one metal ion and one of which contains a mixture of all five metal ions. Obtain a seventh

solution, your unknown solution, from your TA. Using a different capillary tube for each

solution, apply one drop of each of the solutions at the marks on the chromatography paper.

Allow the paper to dry. Then repeat the spotting and drying procedure two more times in order

to increase the metal ion content in each sample spot. N.B. Be careful not to mix up the

capillary tubes! BE PATIENT. It is essential that each spot is dry before applying the

next drop of solution on that spot. If it is not dry, the spots may spread and contaminate

each other.

5. Developing the Chromatogram

Roll the dry chromatography paper into a cylinder with a small gap, ~0.5cm, between

the edges as shown in Figure 2 below. With a piece of tape, ~1.5cm long, connect the two

edges of the paper. Be sure that the two ends of the paper are not touching and that the

cylinder is secure and will not easily come apart. Double check the developing chamber to

ensure that the liquid level is less than 1 cm high. It is essential that when the paper is

put into the chamber, the sample spots are above the liquid level.


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Figure 2.

Remove the plastic wrap from the developing chamber and carefully put the paper

cylinder into the chamber being careful that the paper does not tough the walls of the beaker.

Carefully replace the plastic wrap without disturbing the chamber. Observe what happens as

the solvent moves up the paper and record your observations.

When the solvent has moved up to about 1.5 cm from the top of the paper (~ the 500mL

mark on the beaker), remove the paper from the chamber, unroll it on a clean piece of white

paper and mark the line of the solvent front with a pencil. It is important that this be done

very quickly as the solvent evaporates quite quickly and you must know where the

solvent front was in order to later calculate the R f values. Then recover the developing

chamber with plastic wrap. Allow the solvent to evaporate from the paper and then circle all

of the visible bands with a pencil. Record your observations.

6. Enhancement of the Chromatogram

Roll the paper back into a cylinder, secure the cylinder, and place it in the ammonia chamber

and replace the plastic wrap. When a band for each known cation is visible, after about 5

minutes, record your observations. Remove the paper from the ammonia chamber but leave

it in the fume hood to let the ammonia evaporate. Circle all visible bands. It is important

that this be done quickly as some visible bands disappear. Be sure to tightly cover the


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3. Mark the center of the band for each cation. Measure the distance travelled by each cation.

Calculate the R f value for each cation

4. Calculate the R f value for each of the components of your unknown and by comparison

with the chromatogram of the known ions, identify the components of your unknown.

Writing Assignment

Gas chromatography (GC) is done in the gas phase. Agas chromatograph (GC) has three parts: a sampleintroduction system (injector), an oven containing achromatography column to achieve separation, and adetector. In drug testing, a microliter of liquid urineextract is injected into the injector, a chamber at a hightemperature. The sample is vaporized and swept along ahair-thin glass tube (capillary column, many meters long,flexible enough to be rolled up in a coil) by a carrier gas(mobile phase), such as helium. Different compoundstravel at different speeds because of the differences in

boiling point, polarity, and relative solubility in the carrier gas versus the coating of the inner wall of the column(stationary phase). The compounds emerge from thecolumn outlet at different times after injection (thechromatographic retention time)separated from eachother. Under identical operating conditions, the retentiontime is characteristic of each chemical compound. If twocompounds have the same retention time, they may beidentical (eg, testosterone). If two compounds havedifferent retention times, they certainly are different (eg,testosterone and methyltestosterone). Matching retentiontimes between an unknown and a reference standard isone element of identification. A graph of the amount of substance as a function of the retention time is a

chromatogram.

Fig.. GC-MS data for designer steroid madol .

(A) GC chromatogram; the isomer differs only by the position of the double bond.(B) MS full scan.(C ) MS selected ion monitoring (SIM) scan


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Mass spectrometry (MS) is an analytical chemistry technique used for structure elucidation of unknowns or identification of known compounds. A mass spectrometer has three parts: an ionsource where the compound is ionized to form a molecular ion and fragmented into smaller ions; amass filter that separates ions by mass-to-charge ratio (m/z); and a detector. The graph of ionabundance as a function of m/z is a mass spectrum. In Figure 2B, the molecular ion is 360 andsignificant ions are 345 and 143 (largest = base peak = 100%). The fragmentation pattern isdetermined by weak bonds and other physicochemical characteristics; therefore, fragmentation isreproducible and characteristic of the molecular structure, and the mass spectrum is like afingerprint of the compound. Matching mass spectra between an unknown and a reference standardis another element of identification. Significant ions are so characteristic that matching only threeions (eg, 143, 345, 360) and their percent abundance relative to the most intense of the three (eg,143) has long been widely accepted as proof of chemical identification.

http://www.antidopingresearch.org/BeyondSportsDopingHeadlines.pdf

The combination of gas chromatography and mass spectrometry, GC-MS, is a powerful analytical toolcommonly used in drug analysis. Components of a mixture, commonly a urine sample, are separatedusing gas chromatography and identified from their fragmentation patterns in the mass spectrometer.Describe why the combination of these two techniques is so much more effective an analytical toolthan is either on its own. Imagine yourself to be a lawyer for an athlete who tested positive for ananabolic steroid. In preparing the case for the defense what questions would you ask about the GC-MStesting that was done?


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Name:___________________________ Student No.:____________________

Lab Section No.:___________________

Pre-lab Questions

1. Explain why it is important to mark the chromatography paper with a pencil and not with a

pen.

_________________________________________________________________________

_________________________________________________________________________

_________________________________________________________________________

_________________________________________________________________________

2. What is the stationary phase and what is the mobile phase in this experiment?

_________________________________________________________________________

_________________________________________________________________________

_________________________________________________________________________

_________________________________________________________________________

3. What would happen in this experiment if the solvent level in the developing chamber was more

than 1.5cm high, i.e. if the sample spots were below the solvent level?

_________________________________________________________________________

_________________________________________________________________________

_________________________________________________________________________

_________________________________________________________________________

4. If two cations have the same R f value, how could you detect their presence in a mixture using

paper chromatography?

_________________________________________________________________________

_________________________________________________________________________


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_________________________________________________________________________

_________________________________________________________________________

5. In the chromatogram of a Zn +2 containing solution, the Zn +2 travelled 24mm and the solvent

front travelled 94mm. In another chromatogram of a solution of a mixture of cations, the

solvent front travelled 87mm. If Zn +2 was in that mixture, where would its band appear on the

paper?

_________________________________________________________________________

_________________________________________________________________________

_________________________________________________________________________

_________________________________________________________________________


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Name:____________________________ Student No.:_____________________

Lab Section No.:____________________ Date:__________________________

Data Sheet

Unknown No.:_______________

Room Temperature:___________

OBSERVATIONS

Sample ObservedColour afterDeveloping

Chamber

ObservedColour afterAmmonia

Chamber

ObservedColour withConfirmatory

Reagent

DistanceTravelledby ion (mm)

DistanceTravelledby the

Solvent(mm)

Mn +2

Fe +3

Co+2

Ni +2

Cu +2

5 cation mixture

unknown

unknown


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Experiment 2

BEHAVIOR OF GASES

Part A. DETERMINATION OF THE ATOMIC MASS OF A METALLIC ELEMENT

Introduction

The atomic mass of a metal can be determined by several methods. Two of these are based

upon the reaction of the metal with aqueous acid:

M (s) + n H 3O+ M+n (aq) + n/2 H 2 (g) + n H 2O

The atomic mass can be determined by either (1) determining the number of moles of H 2 gas produced

by a known weight of metal, or (2) determining the number of moles of H 3O+ that are consumed by

a known weight of metal. In this experiment the former method will be used.

Pre-lab Questions

Answer the pre-lab questions on pages 61 and 62. 1-3 refer to Part A of the experiment; 4 and 5, to

Part B..


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Procedure

1. The apparatus for this part of the experiment consists of an 800/1000 mL beaker and a 50 mL

gas burette set up as shown below


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1. Mix 350 mL of distilled water with 150 mL of 1 M HCl solution in an 800/1000 mL beaker.

Fill the gas burette with this solution and, wearing a glove, close the open end with your finger

and invert the burette in the solution remaining in the beaker. Clamp the burette vertically,

with the end immersed just below the surface. Take care that air bubbles are not trapped in the

gas burette. Wash your hands.

3. Obtain an unknown metal from your TA. Clean an approx. 2 cm strip of a 3 mm wide metal

ribbon by rubbing with a small piece of sand paper until the metal is bright and no black spots

are left on the surface. Wipe the clean metal with tissue or filter paper and then weigh it,

accurately to 0.1 mg, without handling.

4. Fold the metal two or three times into a fairly compact mass and press it into a 3-in. test tube.

It should fit snugly against the walls. Fill this tube with distilled water; then, wearing a glove,

insert it, open end up, into the end of the gas burette and lower the gas burette to hold the tube

captive. The acid will diffuse into the test tube and react with the metal. The H 2 evolved will

then collect in the gas burette and displace the dilute acid. Wash your hands.

5. When all the metal has reacted, measure the volume (V, mL) of the gas and the temperature

of the solution in the beaker. It can be assumed that the temperature of the gas in the burette

(T, K) is the same as that of the solution. Also measure the difference in height ( h, mm) of

the solution levels in the burette and in the beaker, and record the atmospheric pressure

(Patm , torr).

Calculations

1. The total pressure of the gas trapped in the burette will differ from the atmospheric pressure

by the hydrostatic pressure exerted by the difference in liquid levels in the burette and in the

beaker. This hydrostatic pressure was measured in millimetres of solution ( h); it can be

readily expressed in torr, since

soln hsoln = HghHg and 1 torr = 1 mm Hg

The density of the solution soln , may be considered equal to that of H 2O at 1.00 g/mL; the


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density of Hg is 13.6 g/mL. Calculate the hydrostatic pressure in torr.

2. The total pressure of the gas trapped in the burette can now be calculated. However, this gas

is a "mixture" of H 2 and H 2O vapour, and

PH2 + P H2O = pressure inside of burette

The vapour pressure of H 2O in the mixture can be estimated from the temperature of the

solution using the data in the Table below, assuming that the vapour pressure of water over the

dilute hydrochloric acid solution does not differ appreciably from that over pure water. Plot

a graph of the data in the Table and read off the P corresponding to the temperature of your

experiment. Note that if the level of liquid in the beaker and the burette are equal,

PH2 + P H2O = P atm .

If the level of liquid in the burette is higher than in the beaker, the pressure in the burette is less

than atmospheric pressure, and

PH2 + P H2O = P atm - P hydrostatic.

Calculate the total pressure of the gas trapped in the burette.

EQUILIBRIUM VAPOUR PRESSURE OF WATER

T ( oC) 16 18 20 22 24 26 28 30

P (torr) 13.6 15.5 17.5 19.8 22.4 25.2 28.3 31.8

3. Calculate the pressure of the trapped H 2 gas.

4. From the values of P H2 , V, and T, calculate the number of moles of H 2 using the ideal gas

relation. Check that your answer here is reasonable as compared to the answer to Pre-lab

Question 3.

5. From the stoichiometry of the reaction and the number of moles of H 2 evolved, calculate the

number of moles of metal in the sample. Assume that n in the balanced equation for the

reaction is either 2 or 3. From the weight of the sample, then determine the two possible


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atomic masses of the metal. Consult the Periodic Table to see which more nearly corresponds

to a known metallic element.

Part B. GRAHAMS LAW OF DIFFUSION

Introduction

Effusion describes the passage of the molecules of a gas through a small hole into an evacuated

chamber. This term is often confused with a similar term, diffusion. Diffusion is the spreading out

of gas molecules through space when a container of gas is opened, allowing the gas to mix freely with

any other gases present.

According to the kinetic-molecular theory of gases, the average velocity of the particles in a

sample of gas is inversely related to the square root of the molar mass of the gas. In the nineteenth

century Thomas Graham, a Scottish chemist, determined experimentally that the relative rate of

diffusion of two different gases at the same temperature was given by the relationship

r 1 /r

2 = (M

2 / M

1)

in which r represents the rate of diffusion of a gas and M its molar mass. This equation, Grahams

Law, is consistent with the kinetic-molecular theory.

While it is not possible experimentally to determine easily and directly the average velocity of

the gas molecules, the rate of diffusion can be determined by measuring how long it takes a gas to pass

through a tube of known length.

In this experiment you will determine the relative rates of diffusion of the gases hydrogen

chloride and ammonia by measuring the distances travelled by the two gases in the same time period.

For a given period of time, the lighter weight gas should diffuse farther than the heavier gas (since

distance travelled in a given time period is directly proportional to rate). The two gases will

simultaneously be introduced to opposite ends of a hollow glass tube. The gases will diffuse through

the tube toward each other and when they meet they will react with each other forming the salt,


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ammonium chloride. The resulting white ring of ammonium chloride that forms will indicate the

position in the tube at which the gases meet.

HCl(g)

+ NH3 (g)

NH4Cl

(s)

Procedure (work in pairs)

CAUTION: This procedure involves the use of concentrated acid and base. Eye protection and

gloves must be worn at all times. The fumes from these chemicals are extremely irritating and

are dangerous to the respiratory tract. These chemicals must be confined to the fume hoods at

all times.

1. Obtain three ~1 cm diameter glass tubes of equal length and two stoppers that will fit snugly

in the ends of the tubes. Measure the length of the tubes.

2. Set up a tube in the fume hood using two adjustable clamps to hold the tube in a steady

horizontal position.

3. In the fume hood will be a small beaker of concentrated HCl and a small beaker of

concentrated NH 3, each covered with a watch glass. (Leave the beakers covered when not in

use.)

4. USING FORCEPS to hold the cotton balls, briefly dip one cotton ball into each of the

solutions. (The cotton balls will dissolve if left too long in the solutions!) Immediately

transfer the cotton balls from the solutions to the opposite ends of the glass tube, inserting the

two cotton balls simultaneously. Stopper the ends of the tube and do not disturb or move the

tube.

5. Patiently watch while the gases diffuse toward each other. When they meet, a white ring of

ammonium chloride will appear. Mark the position on the tube where the ring first appears

and, with a ruler, measure the distance of the white ring from the cotton ball at each end of the


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tube. (As the gases continue to diffuse, the position of the ring will blur so the tube must be

marked as soon as the white ring appears.)

6. Remove the stoppers from the tube and, using forceps , remove the cotton balls and place

them in the large beaker of water that is in the fume hood. Carefully place the glass tube in the

large plastic container in the fume hood.

7. Repeat the experiment with each of the other two tubes.

Calculations

1. From your three sets of data, calculate the average distance diffused by HCl and by NH 3.

2. Using these average distances and assuming that the distance travelled is directly proportional

to the rate of diffusion, calculate the ratio of diffusion rates for the two gases, r NH3 / r HCl .

3. Calculate the ratio of diffusion rates of these two gases that would be predicted by Grahams

Law.

4. Calculate the % deviation of your experimental results from the Grahams Law prediction and

comment on this deviation.


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Name: Student No.:

Lab Section No.:

Pre-lab Questions

1. Would a change in the concentration of HCl used in Part A of this experiment affect the result?

Explain.

2. If H 2SO 4 were used in place of HCl in Part A of this experiment, would this have changed the

volume of gas evolved? Explain.

3. What is the maximum number of moles of H 2 that could be collected in Part A of this

experiment if T = 20 C and Patm = 760 torr?

4. What error would be introduced in Part B of this experiment if the NH 3 was introduced to the

glass tube substantially before the HCl?


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5. How would you expect the result of Part B to differ if the experiment was done at a higher

temperature?


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Data Sheet

Name: Student No.:Lab Section No.: Date:

Name of Lab Partner:______________________________________________________

Part A

weight of metal strip

volume of gas in burette

temperature of solution

h of solution levels

atmospheric pressure

Observations:

Part B

Distance of Ring of NH 4Cl from Cotton in

Trial No. Length of Tube NH 3 End of Tube HCl End of Tube

1

2

3

Temperature

Observations:


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EXPERIMENT 3

Analysis of Antacids by Acid-Base Titration

IntroductionHydrochloric acid is one of the substances found in gastric juices secreted by the lining of

the stomach. HCl is needed by the enzyme pepsin to catalyze the digestion of proteins in the food we

eat. Heartburn is a symptom that results when the stomach produces too much acid (hyperacidity).

Antacids are bases used to neutralize the acid that causes heartburn. Despite the many

commercial brands, almost all antacids act on excess stomach acid by neutralizing it with weak bases.

The most common of these bases are hydroxides, carbonates, or bicarbonates. The following table

contains a list of the active ingredients found in several common commercial antacids, and the

reactions by which these antacids neutralize the HCl in stomach acid.

Compound Formula Chemical Reaction

Aluminum hydroxide Al(OH) 3 Al(OH) 3(s) + 3 HCl(aq) -----> AlCl 3(aq) + 3 H 2O(l)

Calcium carbonate CaCO 3 CaCO 3(s) + 2 HCl(aq) -----> CaCl 2(aq) + H 2O(l) + CO 2 (g)

Magnesium carbonate MgCO 3 MgCO 3(s) + 2 HCl(aq) -----> MgCl 2(aq) + H 2O(l) + CO 2 (g)

Magnesium hydroxide Mg(OH) 2 Mg(OH) 2(s) + 2 HCl(aq) -----> MgCl 2(aq) + 2 H 2O(l)

Sodium bicarbonate NaHCO 3 NaHCO 3(aq) + HCl(aq) -----> NaCl(aq) + H 2O(l) + CO 2 (g)

In this experiment, several brands of antacids will be analyzed to determine the number of

moles of acid neutralized per tablet and the cost analysis of each tablet. The analytical procedure used


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is known as back titration . In this procedure, a known amount of HCl, which is in excess of the base

in the tablet sample, will be reacted with a weighed sample of a ground antacid tablet. The HCl

remaining after the antacid neutralization reaction occurs will be determined by titration with a

standardized NaOH solution to a bromophenol blue endpoint. The number of moles of HCl neutralized

by the antacid is the difference between the moles of HCl initially present in the excess added and the

moles of HCl titrated by the NaOH.. *

*www.chem.latech.edu/~deddy/chem104//104Antacid.htm

Pre-lab Questions

Answer the pre-lab questions on page 71.

Procedure

1. Obtain two different brands of commercial antacid tablets from your TA. For each record the

brand name, the number of tablets in the bottle, the cost per bottle, and the mass of the tablet

and of the components of the tablet as described on the label. Grind one of the antacid tablets

with a mortar and pestle to a fine powder. Weigh out approximately 0.2g of the powder into

each of two 125mL or 250mL Erlenmeyer flasks (the mass should be accurately measured to

0.0001g.

2. Pipette 25.0mL of standardized ~0.1M HCl into each of the flasks and swirl to dissolve the

powder. Be sure to record the actual concentration of the HCl solution.

3. Put the flasks on a hotplate, heat to a very gentle boil with occasional swirling and maintain

the heat for about 1 min to remove any dissolved CO 2.. Remove the flasks from the hotplate

(caution: the flasks are hot) and add about 6 drops of the indicator, bromophenol blue. This

indicator is yellow at pHs below 3.0 and blue at pHs above 4.6. Swirl the flasks. If the

solution turns blue, it means that not enough HCl was added. In this case, pipette an additional
http://www.chem.latech.edu/~deddy/chem104//104Antacid.htmhttp://www.chem.latech.edu/~deddy/chem104//104Antacid.htm


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10.0mL of HCl into the flask and boil again. Repeat these additions of HCl as necessary until

a yellow solution is obtained. Record the total volume of HCl added to each flask.

4. Obtain about 50mL of the standardized NaOH solution in a graduated cylinder. Record the

concentration of the NaOH. If you are the first lab section of the day, rinse the clean burette

with a small portion, ~5mL, of the NaOH solution. Then, using a funnel, fill the burette. Note

that there is no reason to try to fill the burette to exactly the 0.00 mark. Make sure that

there are no air bubbles in the burette tip. Record the initial volume of NaOH (to 2 decimal

places). If you are not the first lab section of the day, there is no need to rinse the burette. Just

fill it and proceed.

5. Make sure that the antacid solution has cooled to room temperature before proceeding. Slowly

and carefully, with constant swirling, titrate the antacid sample with the NaOH solution to a

faint blue endpoint. As the yellow colour disappears, slow down the rate of delivery of the

NaOH to a dropwise flow. When a single drop results in the formation of the blue colour,

stop! Swirl the flask for another 10-15 seconds to make sure the colour remains blue. If it

does not, add one more drop of NaOH and repeat. Read and record the final volume of NaOH

solution on the burette

6. Titrate the contents of the second flask as in 5.

7. Repeat Steps 1 - 6 for the other brand of antacid tablet.

8. Unused antacid powder can be disposed of in a garbage bin. Titrated solutions can be flushed

down the sink with running tap water. If you are in the last lab section of the day, drain any

remaining NaOH that is in the burette into the sink with running tap water. Rinse the burette

twice with tap water, dispensing it through the burette tip, followed by twice with distilled

water. Clamp the burette upside down with the tip open. Rinse each volumetric pipette twice

with tap water and twice with distilled water and clamp them upside down. If you are not in


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the last lab section of the day, leave the remaining NaOH in the burette for the students in the

next lab to use. You do not need to wash the pipette.

Calculations

1. From the concentration of HCl and the total volume of HCl added, calculate the total number

of moles of HCl added.

2. From the concentration of NaOH and the volume of NaOH used in the titration, calculate the

total number of moles of NaOH used in the titration.

3. Subtracting the number of moles of base used in the titration from the total umber of moles of

acid added, gives the number of moles of acid that were neutralized by the basic antacid.

Calculate the number of moles of base in the antacid powder sample. (Dont overlook the

stoichiometry of the reaction!). Then scale this up to calculate the number of moles of base

in the whole antacid tablet. Convert the moles to grams. How closely does this number of

grams agree with the number of grams of active ingredient listed on the label of the bottle?

4. While lay people may think that they should consider the cost per gram of the antacid, chemists

know that cost per mole of base or the cost per mole of acid that can be neutralized, is a more

meaningful measure of cost efficacy. Calculate this cost per mole of acid that can be

neutralized.

5. Think about what factors other than cost per mole of acid neutralized should be considered in

ones choice of an antacid.


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Writing Assignment

This is a role playing experiment. We, the students and staff of CHM110H have been transformed into

the technical staff of GenChem Co., Ltd.*, a quality analytical chemistry testing facility. We have

been contracted by the Ministry of Health to analyze a number of commercial antacids, both for their

therapeutic and their cost efficacy. These data may be used to make decisions regarding the brands

of antacids to be routinely stocked in long-term care facilities. Therefore the writing component of

your report for this experiment is a letter to the Minister summarizing your findings and making your

recommendations. Remember, the Minister is not a chemist so your explanations must be clear in

laypersons terms. Also, the Minister is a busy person, so keep your letter to a single page.

* with permission of Prof. Peter Jeschofnig


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Name:________________________ Student No.:________________

Lab Section No.:________________

Pre-lab Questions

1. In a strong acid-strong base titration at room temperature, the pH at the equivalence point is

7. The indicator bromothymol blue changes colour over the pH range of 3.0-4.5. Explain why

it is okay to use this indicator for this titration.

_________________________________________________________________________

_________________________________________________________________________

_________________________________________________________________________

_________________________________________________________________________

_________________________________________________________________________

2. Why is it important to remove the CO 2 from the solutions before doing the titration?

_________________________________________________________________________

_________________________________________________________________________

_________________________________________________________________________

_________________________________________________________________________

_________________________________________________________________________

3. The analytical technique in this experiment is referred to as back titration. Define this term.

_________________________________________________________________________

_________________________________________________________________________

_________________________________________________________________________

_________________________________________________________________________


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Name:__________________________ Student No.:_____________________

Lab Section No.:__________________ Date:___________________________

Data Sheet

Antacid 1 Antacid 2

Identity, number of tablets per

bottle, cost per bottle, labelinformation

Sample 1 Sample 2 Sample 1 Sample 2

mass of weighing paper (g)

mass of paper + powder (g)

mass of powder (g)

concentration of HCl (M)

total volume of HCl added (mL)

concentration of NaOH (M)

initial burettereading (mL)

final burettereading (mL)

volume of NaOHused (mL)


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EXPERIMENT 4

Spectrophotometric Determination of the Stability Constant of a Complex Ion

Introduction

When a metal ion, acting as a Lewis acid, reacts with a ligand, a molecule or ion with a lone

pair of electrons and acting as a Lewis base, the resulting species is a complex or complex ion if

charged. Some common ligands are H 2O, NH 3, Cl- and heavier halides, CO, CN - and SCN -. Many

transition metal cations form very stable complex ions with many of these ligands and the stability of

these complex ions is expressed by the equilibrium constant for the formation of the complex ion.

Such equilibrium constants are referred to as stability constants or formation constants.

In this experiment an aqueous solution containing ferric ions is reacted with thiocyanate ion

to form a complex ion. The reaction is

[Fe(H 20)6]+3(aq) + SCN -(aq) [Fe(H 20)5(SCN)]

+2(aq) + H 2O

Because the concentration of water in dilute aqueous solution is essentially constant, we normally omit

the waters in the equation and write a simplified version

Fe+3(aq) + SCN -(aq) Fe(SCN) +2(aq)

for which the stability constant is expressed as

K = [Fe(SCN) +2]/[Fe +3][SCN -] .

The object of this experiment is to determine the value of this stability constant.

A knowledge of the factors influencing the stability of complex ions and knowing the value

of their stability constants is of importance in guiding the synthesis of new inorganic materials. As

well, it contributes to our understanding of many biological processes. For example, the fact that the

stability constant for the binding of CO to the Fe +2 of hemoglobin is about 250 times greater than that

for binding O 2, is what is responsible for the toxicity of CO in the atmosphere.


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Pre-lab Questions

Answer the pre-lab questions on page 81-82.

Procedure

Overview: Known concentrations and volumes of the reactants will be mixed and the

resulting concentration of the complex ion product will be measured

spectrophotometrically. From these data, the equilibrium concentrations of reactants

and products can be obtained and the stability constant can be calculated. This is

possible because the product has a characteristic blood-red colour while the reactants

are pale yellow and colourless respectively. Appendix B explains the relationship

between the concentration of a species and its ability to absorb light of a particular

wavelength, the Beer-Lambert Law. Appendix C explains the operation of the

spectrophotometer (colourimeter), the Spectronic 20, that will be used to measure the

absorbance of light.

In Part A of the experiment you will work with a partner to create a calibration

curve for Fe(SCN) +2 , a graph of absorbance of light vs concentration of complex ion.

The solutions used to create this curve all have a very large excess of the SCN -, thus

ensuring that the reaction goes essentially to completion, i.e. that all of the Fe +3 is

turned into the complex Fe(SCN) +2 .

In Part B of the experiment you will work individually. You will prepare five

different mixtures of the reactants, each containing the same amount of Fe +3 but

different amounts of SCN -. In these mixtures the concentrations of the two reactants

will be similar. So the reaction will not go to completion and an equilibrium between

reactants and products will be established. The absorbance of light of each of these

mixtures will be measured and compared with the calibration curve in order to

determine the concentration of the complex ion formed.


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PART A (work with a partner)

1. Turn on the spectrophotometer so that it can warm up while you prepare the solutions for

creating the calibration curve. Label six, clean and dry large test tubes, 0-5. Using the

graduated pipettes provided, prepare the six solutions as described in Table 1. Do not pipette

directly from the stock bottles of reagent. Transfer the required amount of reagent into

a clean beaker and pipette from that. Be careful not to mix up the pipettes. Each pipette

should be used for only one reagent. It is essential that there is no cross contamination

between the reagents. After all of the solutions are prepared, seal each tube with a small

square of Parafilm and thoroughly mix each solution.

Table 1. Composition of the Solutions for Preparing the Calibration Curve

SolutionNumber

0.1M HNO 3(mL)

0.2 M NaSCN (mL)( in 0.1M HNO 3)

0.0005M Fe(NO 3)3 (mL)( in 0.1M HNO 3 )

Total Volume(mL)

0Blank

6 4 0 10.00

1 5.60 4.00 0.4 10.00

2 5.00 4.00 1 10.00

3 4.40 4.00 1.6 10.00

4 3.80 4.00 2.2 10.00

5 3.20 4.00 2.8 10.00

2. Calibrate the spectrophotometer. Using the + or - button on the spectrophotometer, set the

wavelength to 447 nm. Rinse a cuvette with the blank solution, solution 0, and then fill the

cuvette up to the arrow with solution 0. Wipe the outside of the cuvette with a soft tissue to

make sure that it is clean and dry. Place the cuvette into the sample holder of the

spectrophotometer so that the arrow is facing you. Close the cover. Press the mode button


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to activate the abs mode and then press the set reference button to read zero absorbance.

Remove the cuvette and discard the solution in a waste beaker. Use the same cuvette

throughout the rest of the experiment. Once the spectrophotometer has been zeroed, do not

perform any reference setup for the remainder of the experiment. If you accidentally do, repeat

the calibration procedure.

3. Measure the absorbance of the calibration solutions 1 - 5. Rinse the cuvette with solution 1

before filling it to the arrow with the solution. Wipe the outside of the cuvette, place it in the

sample holder, close the cover, and read and record the absorbance. Repeat this procedure for

solutions 2 - 5.

PART B (work individually)

1. Before using the used pipettes, rinse them three times with distilled water and once with the

solution to be used. Label five clean and dry, large test tubes, 6-10, and prepare the test

solutions as described in Table 2 taking the same precautions as in A.1. above. After all of

the solutions are prepared, seal each tube with a small square of Parafilm and thoroughly mix

each solution.

Table 2. Composition of Test Solutions

SolutionNumber

0.1M HNO 3(mL)

0.002 M NaSCN (mL)( in 0.1M HNO 3)

0.002M Fe(NO 3)3 (mL)( in 0.1M HNO 3 )

Total Volume(mL)

6 4.00 1.00 5.00 10.00

7 3.00 2.00 5.00 10.00

8 2.00 3.00 5.00 10.00

9 1.00 4.00 5.00 10.00

10 0 5.00 5.00 10.00


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2. Rinse the same cuvette that was used in Part A, first with distilled water and then with solution

6. Then fill the cuvette up to the arrow with solution 6. Wipe the outside of the cuvette with a

soft tissue to make sure that it is clean and dry, place it in the sample holder so that the arrow is

facing you, close the cover and read and record the absorbance. Repeat this procedure for

solutions 7 - 10.

3. Dispose of all solutions from Parts A and B in the waste bottles in the common fumehoods.

Remove the labelling tape from all test tubes used. Rinse the test tubes with tap water and put

them in the used test tube bin. Rinse the pipets and the cuvette twice with tap water and twice

with distilled water. Ensure that your lab bench and fume hood are clean and dry.

Treatment of Data

PART A

1. Calculate the concentration of Fe +3 put into each of the solutions, 1-5. Since SCN - is in large

excess in each of the solutions, Fe +3 is the limiting reagent. So the concentration of the Fe(SCN) +2

complex formed is equal to the concentration of Fe +3 put into the solution. (See Pre-lab Question

1).

2. Plot a graph of absorbance vs concentration of Fe(SCN) +2 . This is your calibration curve. You

will use it to determine the concentration of Fe(SCN) +2 in each of your test solutions in Part B.

PART B

1. Calculate the moles of Fe +3 put into each solution, 6-10, and the initial concentration of Fe +3 in

mol/L.


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2. From the measured absorbance for each solution, read the equilibrium concentration of Fe(SCN) +2

off the calibration curve.

3. Subtracting the concentration in 2. from that in 1. gives the concentration of Fe +3 remaining at

equilibrium.

4. Calculate the moles of SCN put into each solution, 6-10, and the initial concentration of SCN

in mol/L.

5. Subtracting the concentration in 2. from that in 4. gives the concentration of SCN - remaining at

equilibrium.

6. From the equilibrium concentrations for each solution, 6-10, calculate the stability constant for

the reaction.

7. Calculate the average of the stability constants determined.

Writing Assignment

Assume that the spectrophotometers in the first-year chemistry lab have reache

Documents

Course Manual 13-14